Test for Topics Chemistry Form 3

TOPIC : 1 CHEMICAL EQUATIONS

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

CHEMICAL EQUATIONS

NAME………………………………………..CLASS………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, and B

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

SECTION A. 20 MARKS.

1. MULTIPLE CHOICE QUESTIONS

he equation X2+ + 2e- X, represent:-

Oxidation reaction

Reduction reaction

Displacement reaction

Synthesis reaction

Neutralization

The reaction between Silver nitrate and Sodium chloride to form Silver chloride and Sodium nitrate is an example of a ………………. Reaction.

Direct combination

Simple displacement

Double displacement

Decomposition.

Ammonium chloride reacts with sodium hydroxide solution on warming. The net ionic equation for the reaction is.

H+(aq) + OH-(aq) → H2O(l)

NH+(aq) + OH-(aq) → NH3(aq) + H2O(aq)

Na+(aq) + Cl-(aq) → NaCl(aq)

2NH4+(aq) + 2Cl-(aq) → NaCl(aq) + Cl(g) + H2(g)

The percentage composition of compound K is 53.3% oxygen, 6.7%hydrogen and 40%carbon by mass.The empirical formula is;

C2H4O

CH2O

CH4O

C2H2O

If 1g of hydrogen is exploded in air, the mass of water formed is:

1.8g

18g

The ionic equation of the reaction between hydrochloric Acid and Sodium hydroxide is;

a+ + Cl- NaCl

a+ + OH- NaOH-

3O+ + Cl HCl + H2O

3O + Cl HCl + H2O

H+ + OH- H2O

Write the letter of the best match from column B against a statement in column A.

LIST A

LIST B

A mechanical equation

A decomposition reaction

A synthetic reaction

Double displacement

Neutralization reaction.

Spectator ions

A reaction that proceeds in both directions

A well written balanced chemical equation

Decomposition of a substance using heat

Writing equations using ions and not molecules

Fe(NO3)2(aq) + Cu(s) → Cu(NO3)2(aq) + Fe(s)

→ Fe(s)

Donot take part in reactions

Are cancelled in writing equations

Reversible reaction

Irreversible reaction

Cu(s) → Cu

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

H2(g) + O2(g) → H2O(g)

Thermomal decomposition

Heat energy

Ionic equation

Molecular equation

H2 O2 + O2(l) MnO2 H2O(l) + O2(g)

BaCl2(aq) + H2SO4(aq) → BaSO4(s) + 2HCl(aq)

CaCO3(s) + 2HNO3(aq) →Ca(NO)2(aq) H2O(l) + CO3(g)

Ag+(aq) + Cl-(aq) → AgCl(s)

Cu(s) + 2HNO3(aq) → N.R

SECTION B.

Answer the questions in this section in brief

Rewrite complete and balance the following chemical equation of reactions:

n(s) + Cl2(g)

aO(s) + H2O(l)

+ N2(g) Li3 N(s)

2CO3 + HCl(aq)

b(NO3)2(aq) + Na2SO4(aq)

Write ionic equations from the following chemical reactions:

Magnesium reacts with dilute sulphuric acid

Silver nitrate solution reacts with iron (II) chloride solution

Dilute solution acid reacts with solid calcium carbonate

Iron displaces copper of copper (II) sulphate solution

Chlorine gas displaces iodine of potassium iodide solution

(a) (i) Define

Ionic equation

Molecular equation.

(b) Complete and balance the following chemical equations.

(i) KCLO3

heat

b (N03)2

2 S04 + KOH

UCO3

a + H2O

With the aid of a well balance chemical equations explain what would happen in each of the following reactions

Potassium carbonate is strong heated

Concentrated sulphuric acid added to dry salt and heated

Sodium nitrate heated

Concentrated sulphuric acid slowly acted to the sugar

Soluble alkali added to the soluble salts

Carbon dioxide passed through lime water

Water is added to a white copper (11) sulphate

A glowing splint of wood is lowered into a jar of (i) Hydrogen gas (ii) Carbon dioxide gas

Ammonium chloride is heated

Dilute nitric acid was heated with (i) warm copper oxide (ii) Zinc carbonate

Write the product and balance the following chemical equations

g N03(aq) + NaCl(aq)

n(s) + H2S04(aq)

gCl2(aq) + AgN03(aq)

a2S04(aq) + BaCl2(aq)

nC03(s) + HCl(aq)

a(s) + H20(l)

Complete and balance the following chemical equations:-

e(OH)3 + H2 SO4

gCO3 + HCl

a + O2

3PO3 + Al2O3

Explain the meaning of each of the following types of chemical reactions and support your explanation with the help of a relevant chemical equation.

Synthesis (combination) reaction

Decomposition reaction

Precipitation reaction

Single displacement reaction

Neutralization reaction

(a) Define:-

Redox reaction

Oxidation reaction

Reduction reaction

(b) Classify the following reactions into oxidation and reduction reactions:-

(s) + O2(g) SO2(g)

2(g) + 3H2(g) 2NH3(g)

e2+ (aq) – e- Fe3+(aq)

e3(aq) + e- Fe2+(aq)

(

i) Hydrogen + chlorine Hydrogen chloride

(

ii) Magnesium oxide + carbon Magnesium + carbon monoxide

(iii)Silver nitrate + sodium chloride silver chloride + sodium nitrate

(a) Explain the meaning of “an ionic equation”

(b) Write an ionic equation for each of the following:

Carbon dioxide gas dissolves in an aqueous solution of potassium hydroxide

The reaction between magnesium and dilute hydrochloric acid.

School Base-Online Page 4

TOPIC : 2 HARDNESS OF WATER

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

HARDNESS OF WATER

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

SECTION A 20 MARKS.

Hard water which can be softened by boiling method contains dissolved:-

Calcium Hydrogen sulphate

Magnesium chloride

Magnesium carbonate

Calcium sulphate.

Which of the following compounds can be used to remove hardness of water?

Calcium chloride

Sodium carbonate

Magnesium hydroxide

Potassium hydroxide

Zinc chloride.

……………….. Is a one of the substance whose dissolution in water caused permanent hardness?

Carbon dioxide

Calcium sulphates

Magnesium carbonate

Sodium carbonate

One of the disadvantages of hard water is that is;

Causes corrosion of water pipes

Causes increased tooth decay

Requires more soap for washing

Contains minerals that are harmful

Hard water which is softened just by boiling contains dissolved;

Calcium carbonate

Calcium chloride

Sodium carbonate

Magnesium sulphate

Calcium hydrogen carbonate

Which of the following methods can be used to remove both temporary and permanent hardness of water.

Using ion exchange chamber

Boiling

Using of calcium hydroxide

Evaporation

Which of the following is not an advantage of hard water?

Suitable for drinking

Strengthens bones and teeth

Suitable for brewing of beer

Makes our teeth coloured

One advantage of soft water is that;

Forms lather with soap

Suitable for drinking

Has a nice taste

Used in treatment of ulcers

Which is the best method that removes hardness of water to a large extend?

Boiling

Distillation

Using washing soda

Using calcium hydroxide

The main source of salt water on earth is;

Sea

Springs

Glaciers

Wells

2. Matching Items Questions.

LIST A

LIST B

Water that easily forms lather with soap

Water that does not form lather with soap

Water that contains dissolved calcium and magnesium hydrogen carbonate

Water that contains dissolved sulphates of calcium and magnesium

An element whose complex ion is used in softening water

Substance formed when soap is used in washing using hard water

An element which forms constituent of eggs.

Substance formed on kettles when used to boil hard water.

A gas that dissolves in water in soils to cause hardness

Another name for hydrated calcium sulphate

Plaster of Paris

Gypsum

Nitrogen dioxide

Carbondioxide

Calcium

Phosphorous

Soft water

Hard water

Permanent hard water

Temporary hard water

Scum

Stain

Fur

Coating

Sodium

Potassium

magnesium

Write down the chemical equations used when softening water of the

Temporary hardness through (i) boiling water (one question) (ii) use of chemicals (two equations)

Permanent hardness through (i) use of chemicals (one equation) (ii) iron exchange (one equation)

Define the following terms;

Soft water

Hard water

Permanent hardness of water

Temporary hardness of water

a) What is the hardness of water?

b) Briefly explain types of hard of water.

c) State the causes of hardness of water for each type mention in (b) above.

d) Explain how you would remove the hardness of water according to its type.

e) Give three (3) advantages and three (3) disadvantages of the hard water.

Soap solution of different amount of water are tested from four different sources and produces lather observed for 30 seconds. The three groups of waster were the untreated; boiled and treated by ion exchange. The results are as follows;

Sample	Volume of soap (cm3) used for water that was
Sample	Untreated	Boiled	Passed through ion exchanger
A B C D	12 17 26 1.6	1.8 17 20 1.6	1.8 1.7 1.8 1.6

Use the above results to answer the following questions;

Which is the hardest water sample? Why?

Which sample is like distilled water? Explain.

Which chemical substance might be the cause of hardness in (i) Sample A (ii) Sample B?

Write an equation for the reaction of removing hardness in sample C.

(a) State two advantages of hard water.

(b) State two disadvantages of hard water.

(a) Distinguish between;

Acidic salt and basic salt

Hydrated and anhydrous salt

(b) Vinegar, lemon juice and yoghurt all have sour taste. Give the other possible properties would you expect them to have? Give 3 points.

(a) Give meaning of the following;

Basicity of an acid

The pH scale

(b) Name substances which when dissolved in water causes;

i. Temporary hardness of water

ii. Permanent hardness of water

i. Temporary hardness of water can be removed by boiling. (3 marks)

ii. Permanent hardness of water can be removed by chemical means.

Give a brief account to the following

Why does water note have any effect on litmus paper?

(i) What would happen to a well stoppered bottle full of water left in a deep freezer over the night? Why does this happen?

(ii) Why is iron not usually recommended in the construction of steam pipes and boilers.

(i) Name two compounds which may cause temporary hardness of water and two

ions which largely cause permanent hardness of water.

(ii) Write one balanced chemical equation which shows how temporary hardness can be removed by boiling. Also write one balanced equation which shows how sodium carbonate can be used to remove permanent hardness of water.

(a) What do you understand by the following terms:-

Abase (ii) Alkali (iii) Weak acid (iv) Hygroscopic substance.( 2 marks)

(b) With the aid of chemical equation describe how permanent and temporary hardness of water can be removed?

(a) State two advantages of hard water.

(b) State two disadvantages of hard water.

School Base-Online Page 5

TOPIC : 3 ACIDS, BASES AND SALTS

School Base-online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

ACID BASES AND SALTS

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

1. MULTIPLE CHOICE QUESTIONS

Sea water contains various salts. Which salt is present in the largest proportion?

Magnesium sulphate

Sodium chloride

Calcium sulphate

Magnesium chloride

Which pair among the following comprises of good drying agents?

Calcium carbonate and lead (II) chloride

Anhydrous calcium chloride and concentrated sulphuric acid

Calcium sulphate and calcium oxide

Concentrated nitric acid and concentrated hydrochloric acid

A weak acid is the best describe as ……………………..

An acid that does not ionize completely.

A dilute acid.

An acid that does not react with any substance.

An acid that is harmless.

When dilute solutions of calcium chloride and sodium carbonate are mixed;

A white precipitate of sodium chloride is formed

A white precipitates of calcium carbonate is formed

A colourless solution of calcium carbonate and sodium chloride are formed

A mixture of precipitates of sodium chloride and calcium carbonate are formed.

Insoluble salts like lead (II) chloride, generally can be obtained in the laboratory by;

Crystallization

Precipitation

Decomposition

Evaporation of its concentrated solution

A hygroscopic substance;

Evaporates to the air

Loses water to the air

Add water to a compound

Absorbs moisture from the air

Removes elements of water from a compound

An efflorescent substance;

Evaporates to the air

Loses water to the air

Absorbs moisture from the air

Removes elements of water from a compound

Adds water to a compound

Which of the following is a weak acid?

Sulphuric acid

Hydrochloric acid

Ethanoic acid

Nitric acid

A solution of a PH of 1 is said to be;

Slightly acidic

Slightly basic

Strong acid

Neutral

Which of the following is not a product when silver nitrate is heated?

Silver oxide

Silver metal

Nitrogen dioxide gas

Oxygen gas

2. Matching items questions.

LIST A

LIST B

A salt that absorbs water to form solution

A salt that loses water into the atmosphere

A salt that absorbs water from atmosphere but do not form solution

Number of replaceable hydrogen ions

An acid whose basicity is three

A salt that is used in softening water

A metal nitrate which when heated decomposes to form metal nitrite and oxygen

A gas with a brown colour

An oxide which is orange when hot and yellow when cold

Reaction between an acid and a base to form salt and water only

Nitrogen I Oxide

Nitrogen dioxide

Lead oxide

Zinc oxide

Basicity of an acid

Delinquent

Hygroscopic substance

Efflorescent

Potassium nitrate silver nitrate

Hydrating salt

Drying agent

Calcium sulphate

Sodium carbonate

Neutralization reaction

Sulphuric acid

Phosphoric acid

Displacement reaction

A) Acids can react with, carbonates, hydrogen carbonates, oxide and hydroxide. Presence of carbon dioxide can be proved using lime water. Support the above statements by using balanced chemical equations.

B) How can one prepare a flower extract indicator and use it?

Distinguish between the following terms

Hydrate salts and unhydrate salts

Efflorescence and deliquescence

Base and alkali

Acid and base

Acid salts and normal salts

Fill the following tables

Table one.

Test	Observation
1. Blue and red litmus paper dipped into sodium hydroxide
2. A piece of magnesium metal dropped into diluted hydrochloric acid
3. One spatula of sodium carbonate added into dilute sulphuric acid
4. 2cm3 of AgN03 solution added to the solution of mgcl2

Table Two.

Indicator	Colour change in acidic solution	Colour change in alkaline solution
1. Phenolphalein
2. Methyl orange
3. Litmus paper

(a) 416g of anhydrous barium chloride where obtained when 488g of hydrated salt were heated.

Calculate the value of n is the formula BaCl2.nH2O (Ba = 137)

(b) Give the meaning of the following;

(i) An acid

(ii) Molar solution of an acid

(iii) Give three characteristics of acid.

Differentiate salts behave differently when heated. Complete the following equations that show the actions of heat on certain salts;

PbCO3(s)

2Ca (NO3)2(s)

(NH4)2 CO3(s)

2NH3(g)+

NH4Cl(s)

(a) Distinguish between strong acid and weak acid

(b) Explain three applications of neutralization

25cm3 of potassium hydroxide were placed in a flask and a few drops of phenolphthalein indicator were added. Dilute hydrochloric acid was added until the indicator changed colour. It was found in the 21cm3 of acid were used.

From above information answer the following questions;

A: (i) What piece of apparatus should be used to measure out accurately 25cm3 of sodium hydroxide solution?

(ii) What color was the solution in the flask at the start of the titration?

(iii) What color did it turn when the alkali had been neutralized?

B: (i) Was the acid more concentrated or less concentrated than the alkali?

(ii) Name the salt formed in the neutralization.

(iii) Write an equation for the reaction.

(iv) Is the salt, normal or acidic salt? Give reasons for your answer.

C: (i) Utilizing the given information of question 9 (a) and (b) above describe how you can obtain pure crystals of the salt.

(ii) Is it a soluble salt or insoluble salt? Give a reason for your answer.

A sample of water forms scum with soap. When the water was boiled and then cooled it still formed scum.

Does this water have temporary hardness or permanent hardness?

Suggest one treatment that can be used to soften the water.

(a) Define the term indicator.

(b) Write only two uses of pH scale.

(a) Give meaning of the following;

Basicity of an acid

The pH scale

(b) Name substances which when dissolved in water causes;

i. Temporary hardness of water

ii. Permanent hardness of water

i. Temporary hardness of water can be removed by boiling.

ii. Permanent hardness of water can be removed by chemical means.

(a) Differentiate between basic salt and acidic salt.

(b) Categorize the following salts;

i. PbSO4 ii. MgHSO4 iii.Zn(OH)Cl iv. MgHCO3

v. NH4HSO4 vi. Ba(NO3)2

Salts	Category

(a) Define the term “neutralization reaction” (give one example)

(b) Write down the names and formulae of three common acids in the laboratory.

(c ) What is an indicator? Give four (4) examples of acid- base indicators.

(d) Write down the products formed when each of following pairs of compounds react:

(i) Acid and metal

(ii) Acid and metal carbonate

(a) Name of element found in all acids

(b) Describe two tests you would carry out on a liquid to decide whether it is an acid or not in each case describe the test and result you would expect.

(c)If some acids spills on your clothes, which of the following substance would you put on the cloth to prevent further damage? Sodium hydroxide solution, sodium hydroxide pellets, concentrated ammonia solution or sodium hydrogen carbonate? Why?

School Base-online Page 6

TOPIC : 3 ACIDS, BASES AND SALTS

P CLASS="western" ALIGN=CENTER STYLE="margin-bottom: 0.11in">CHAPTER 3

ACIDS, BASES AND SALTS.

ACID AND BASES.

KEY TERMS AND CONCEPTS.

Acid- a substance that dissociates in water to produce hydrogen ions as the only positive ions.

Base – this is a substance that ionizes in water to form hydroxyl ions as the only negative ions.

Strong acid- this is an acid that ionizes completely when dissolved in water.

Weak base/acid- are acids or bases that ionizes partially when dissolved in water

Salt – an ionic compound formed when a cation derived from a base combines with anion derived from an acid.

Complex ions- is formed when simple metallic cations joins with two or more anions to form a large ion.

Delinquent salt- this is salt that absorbs water from the atmosphere forming a solution.

Hygroscopic salt- this is a salt that absorbs water from the atmosphere but does not form a solution.

Efflorescent- this is a salt that looses water into the atmosphere becoming unhydrous

Basicity of an acid- this is the number of hydrogen ions that an acids contains and can ionize in water.

An indicator- this is a substance that is used to test for the presence of an acid or base on a substance.

Neutralization- this is the reaction between an acid and a base to form salt and water only.

Hydrated salt- this is a salt that contains water of crystallization.

Natural sources of acids and bases.

Acids occur naturally in sour substances. These substances include citrus, fruits like oranges, lime and sour milk.

Acids of are also found in manufacturing of substances like

Banana peels

Baking powder

Tooth paste (manufactured)

Wood ash

Soaps

Bleaches

Cleaning agents.

Acid have sour tastes and turns blue litmus paper to red.

Bases have a bitter tastes and turns red litmus paper to blue.

NB: is a special dye that can be used either as a solution or on paper.

Some common natural acids and their sources

Name	Sources
1. carbon acid 2. oxalic acid 3. formic 4. Acetic acid 5.lactic acid 6.citric acid	1. soft drink 2. spinach 3.Ants 4. vinegar 5.sour milk 6. citrus fruits

Table 3.1

Some common natural bases and their sources

Name	sources
1. Ammonia 2. methylamine 3. ethylamine 4. pyridine 5.putresure 6. cadaverine	1. gills of fish 2.decaying fish 3. decaying fish 4. coal tar 5. decaying meat 6. decaying flesh

Table 3.2

Laboratory acids and bases.

Acids

hydrochloric acid (HCL)

Sulphuric acids (H2SO4)

Nitric acid (HNO3)

Bases

Ammonium solution

Metal oxides

Metal hydroxides

ACID: Is a chemical substance with when dissolved in water produces (H+) hydrogen ions as charged ion he only positively charged ions.

When a acids react with some substances the hydrogen in the acid is displaces and substances are formed.

Physical properties

Most of dilute acids have a sour taste

Acids turn blue litmus paper to red

Most acids are corrosive

Chemical properties

Reaction with metals

Some metals react with dilute acids hydrochloric or sulphuric acids to liberate hydrogen gas.

These metals include Zinc, Aluminum and Magnesium.

Caution: Metals such s sodium and potassium should not be used in such experiments. This is because the reacts vigorously with an acids. Dilute Nitric acid with metals hydrogen gas is not liberated, because nitric acid is very strong oxidizing agent. It react with metals to give nitrogen (IV) Oxide. However dilute nitric Acid react with Magnesium oxide.

Reaction with bases

Base is an oxide or hydroxide of metals.

Acids react with bases to produce salt and water only.

aO(S)+ 2HCl(aq) CaCl2(aq) + H2O(aq)

Neutralization reaction

NaOH(aq) + H2SO4 Na2SO4(aq) + Na2SO4 + 2H2O(l)

Action of acid with carbonates and hydrogen carbonates.

aCO3(s) + HCl (aq) CaCl(aq) + H2O(s)

a(HCO3)2 +H2SO4 CaSO4 + H2O(l) +CO2(g)

Ionic reaction

a2+ + CO32- + 2H+2Cl- Ca2+ +2Cl-+ H2O(l) + CO2(g)

H++ CO32 H2O +CO2

BASES

Is a substance that produces hydroxyl ions (OH-) which are negatively charged ions.

aOH Na+ + OH-

A base can also be defined as metal oxidized as metal oxide or hydroxide.

ALKALIS

Is a soluble base.

CLASSIFICATIOB OF BASES

DIAGRAM FOR CLASSIFICATION OF BASES.

NEUTRALIZATION

Is a reaction where acids and bases react to produce salt and water only

PROPERTIES OF BASES

Physical properties

Most bases have bitter tastes.

Base turns red litmus paper to blue.

Bases have a (soap) "soapy” or slippery feel.

Bases are oxides and hydroxides of metals and most are insoluble in water.

Soluble gases are alkalis e.g. NaOH, KOH and NH4OH.

Insoluble bases are include e.g. MgO, Cu(OH)2, Al2 (OH)3

Chemical Properties

Bases react with acids to form salt and water.

OH + HCl KCl+ H2O

Alkalis (soluble bases) precipitate insoluble metals hydroxides from their salt solution

3KOH(aq)+ FeCL3(aq) Fe(OH)3(S) + 3KCl(aq)

Alkalis react with ammonium salts to produce ammonium salts to produce ammonia gas

2NH4Cl(s)+Ca(OH)2 CaCl2(aq) + 2H2O(S) + 2NH3(g)

INDICATORS.

Are compounds that show a definite change of color or color change when in an acid or base

They are used to test whether the substance is acidic, basic or neutral.

Examples:

litmus papers

phenolphthalein (P.O.P)

Methylorange (M.O)

Bromothymol blue.

Color change is common indicators when in acid and basic solution.

Indicator

Litmus solution

Phenolphthalein

Methyl orange

Bromothymol blue

Acid

Red

Colorless

Red

yellow

Basic

Blue

Red

Yellow

Blue

Table 3.3

PH Scale

IS the scale related to the concentration of hydrogen ions in a solution.

The higher the hydrogen (ions) (H+) in a solution.

The higher the hydrogen ions in a solution

The higher the hydrogen ions (H+) Concentration, the lower the PH Value.

Figure 3.1

Acid solutions have PH values that are less than seven. The smaller PH value, the more acidic solution is.

Neutral liquid e.g. (H2O) have a PH value, the more acidic solution is.

Neutral liquid e.g. (H2O) have a PH of seven(7).

The solution which PH value of greater than seven are alkaline. The greater the PH the more alkaline a solution is.

UNIVERSAL INDICATORS.

Is the mixture of indicators?

It displays arrange of colors in solutions depending on how acid or basic a substance is.

It is a composed by Ph chart showing the range of colors with red on extreme left (most acid) and blue on the extreme right (most basic)

STRENGTHS OF ACIDS AND ALKALIS.

Acids and alkalis can be classified as strong or weak. This depends on the degree to which they dissociate when dissolve in water.

Strong acids.

These are the ones which dissociate completely in water to give hydrogen ions and he negative ions associated with the alkalis.

2SO4 + H2O H+ +HSO4-

String alkalis dissociate completely in water to give hydroxyl ions and the positive ions associated with the alkalis

aOH H2O Na+ + OH-

Weak acids and alkalis only dissociate partially in water

Strong acid will have PH value between 0 and 2.

Strong alkalis have PH Value between 11 and 14.

Strong and weak acids

Strong acids	Weak acids
Hydrochloric acid Nitric aid Sulphiric acid	Ethanol acid Carbonic acid Methanoic acids

Table 3.4

Concentration of Acids And Alkalis.

The number of hydrogen atoms per molecule of acid can be displaced by a metal in solution.

e.g. hydrochloric acid (HCL) which as an hydrogen atoms that can be displaced by a metal in solution

sulphuric acid (HCl) which has an hydrogen atoms that can be displaces is said to be monobasic acid

sulphiric acid (H2SO4) Has a basicity of two and said to be dibasic acid.

Phosphoric acid (H3PO4) has a basicity of three and said to be tribasic acid.

ACID	BASICITY
HCl H+ + Cl- H2SO4 2H+ +SO2-4 H3PO4 3H+ + PO43- CH3COOH H+ + OH-3 +CO	Monobasic Dibasic Tribasic monobasic

Table 3.5

Neutralization

Is the reaction which involves acids and basic to produce salt and water as the only products.

The reaction is said to be neutral because it produces salt and water that have a PH of 7.

O + HA MA + H2O

ase + Acid Salt + water

Is the same as double displacement reaction

The value at acid that has neutralized the alkali is called a titre.

The method used as called volumetric titration

Application of Neutralization

Treating insect stings such as those bees can inject an acid liquid into the skin can be neutralized by rubbing baking soda (sodium hydrogen carbonate) on the affected area which is a base.

Ant stings or bites contain methanoic acid. They can also be neutralized by using banking soda or any other alkaline substance like cucumber or avocado.

Also wasp stings are alkaline substance and they can be neutralized by acetic found in vinegar.

Soil treatment

Most plants grow well in soils that are neither too acid nor basic. This means the best soil condition is Ph of 7.

When soils are too acid chemicals such as lime Ca(OH)2 are added to adjust the soil PH.

Treating factory wastes.

Liquid wastes from the factories after contain acid. If the waste gets into the river ( waste bodies), it could kill fish and other forms of life. This is usually prevented by adding slaked lime (calcium hydroxide) to neutralize it.

Manufacturing of fertilizer

Ammonia fertilizers are produced by neutralization. A good example is ammonium nitrate which is formed by the action of ammonia gas and nitric acid.

.g. 1.NH3 (g) + HNO3(aq) NH4NO3 (aq)

. NH4OH + HNO3 NH4NO3 +H2O

Base + nitric acid fertilizer + water.

Reducing acid rain occurrence.

Acid rain is caused by chemical reaction that begins when compounds like sulphur dioxide and nitrogen dioxide are released in air. These compounds like react with rain water to form acid rain increases acidity of soils, rivers and lakes and adversity affects vegetation and aquatic organism. To reduce this problem air pollution devices are fitted in exhaust pipes and chimneys to neutralize the acidic compounds before reaching the atmosphere.

Reliving indigestion

Indigestion refers to pain or discomfort in the stomach that associated with difficulty in digestion of food it is caused by the process of excess acid (dilute hydrochloric acid) in the stomach.

Contain magnesium or sodium bicarbonate (antacids)

Neutralizing accidental spills.

If an acid or an alkali, splls on the floor or work surface in the laboratory it can be neutralized e.g .concentrated Rated sodium hydroxide (corrosive ) could be neutralized by adding hydrochloric acid to it.

USES OF ACIDS.

Acids have the following uses.

They have been usd in neutralization process.

Acids are also used as reagents in laboratories

Uses sulphuric acid in removing rust from metals.

Used in manufacture of synthetic textile such as nylon.

In making fertilizers

In production of sulphates

Uses of hydrochloric acid

Production of chloride

In electroplating

Used as catalyst and solvent in organic synthesis.

In manufacture of fertilizers, textiles and runner.

Uses of nitric acid.

Manufacture of fertilizers, explosives and rocket propellants.

Manufacture of nylon

Testing for genuine precious stones

As an oxidizing agent to clean metals

Synthesis of several dyes.

Uses of phosphoric acid.

Manufacture of fertilizers, explosive and rocket propellants.

Manufacture of nylon

Testing for genuine precious stones.

As an oxidizing agent to clean metals.

Synthesis of several dyes

Uses of phosphoric acid.

Manufacture of phosphates for softening water, detergents and fertilizers.

Rust proofing metals

Ingredients in soft drinks and dental cements

As acatalyst.

USES OF ALKALIS

Alkalis can be used in

In various neutralization process

In manufacture of cleaning agents

In various industrial process

e.g. sodium hydroxide is used in the manufacture of paper soap and ceramic.

SALTS.

Is compound formed when the hydrogen ion of acid is replaced directly or indirectly by the metal or ammonium ion.

Is a compound formed other than water when on acid neutralize a base.

Natural Sources of Salts.

Salt can be found in seawater and some rocks.

Such salts are sodium chloride (NaCl), sodium Sulphate (Na2SO4) and calcium carbonate (CaCO3)

Preparation of Salts.

There are several ways of preparing salts the methods depends whether the salt is soluble or insoluble.

Preparation of Soluble Salts

These are several ways of preparing salts. The method depends on whether the salt is soluble or insoluble.

Preparation of soluble salts.

Reaction of acids with metals

Zinc, iron and magnesium can be used to prepare salts.

etals + Acid salt + hydrogen

g + 2HCl MgCl2 + H2

e + H2SO4 FeSO4 + H2

Crystallization

Is the process of formation of solid crystals from homogenous solution.

Crystallization point.

Is the temperature at which crystals form as a solution is being evaporated. Reaction of acid with alkalis.

This is the best preparation used for salts of Na, K and NH3

.g 2NaOH + H2SO4 NaSO4 + 2H2O

H4OH + HCl NH4Cl + H2O

Reaction of acids with insoluble gases

cid + insoluble base Salt + water

UO + 2HCl CUCl2 + H2O

nO+ H2SO4 ZnSO4 + H2O

Reacting acids with metal hydrogen carbonates

CaCO3, CuCO3 and K2CO3 are used.

aCO3 + 2HCL CaCl2 + H2O + CO2

uCO3 + H2SO4 CuCl2 + H2O

2CO3 +2HCL 2KCL + H2O + O2

Reacting acids metals hydrogen carbonates.

aHCO3 + HCl NaCl +H2O +CO2

Direct combination

If the salt required is anhydrous ( without water) form they can be prepared by the direct combination or synthesis.

Preparation of iron III Chloride

Dry chloride is passed over a hot iron wire or wool. The iron III Chloride vapor produced is condensed to give a sold salt.

Figure 3.2

Preparation of insoluble salt

The insoluble salts cannot be prepared by the any of the last methods above.

The salts can be prepared by the method known as ionic preparation or double decomposition.

In this method two salts are used to form insoluble salt and soluble salts,

The insoluble salts precipitate while insoluble remains in a solution. The insoluble remain in a solution. The precipitate is filtered and washed with distilled water and dries.

.g. .Pb (NO3)2(aq)+ 2KI (aq) PbI3(s) + 2K(aq) +NO3(aq)

.g. Pb2+ + 2I- PbI (s)

Other common insoluble salts which can be prepared by using this method include (CaSO4)

Calcium sulphate (CaSO4)

Magnesium carbonate (MgCO3)

Silver chloride (AgCl)

Barium Sulphate (BaSO4)

Barium carbonate (BaCO3)

Lead II Sulphate( PbSO4)

.g. BaCl2(aq) + H2SO4 BaSO4(s) + 2HCL

gNO3(aq) + BaCl2 Ba(NO3)2 + AgCl (s)

Types of Salts.

These are four main types of salts

Normal salts

Is the one which will all replaceable hydrogen ions of an acid are replaced by metal. Example potassium (K2CO3), sodium sulphate (Na2SO4), Ammonium Sulphate (NH4)SO4) , Zinc chloride (ZnCl2)

Acidic salts.

Are formed whwen only part of the replaceable hydrogen ions molecule of an acid is replaced. Example NaHCO3, NaHSO4, KHCO3, KHSO4, NaHPO4

Basic salts.

Are formed when the amount of acid requires neutralizing an alkali or base is insufficient. Examples

Basic Zinc chloride (ZnCl2. Zn (OH)

Basic magnesium chloride (MgCl2, MgOH)

Basic lead carbonate (PbCO3.Pb (OH) 2)

Basic calcium sulphate (CaSO4. Ca (OH)2

Double salts

Are formed when the solutions of two cells salts are mixed up together and then allowed to stand for some time, they react to form a single new salt.

The new salt is different from the original salts.

.g. (NH4)2SO4 + Fe2(SO4)3 FeNH4(SO4)2.12H2O

Iron (III) ammonium sulphate (Iron alum)

Solubility of Salts.

Is the number of grams of the solute required to saturate 100g of solvent at a given temperature?

Solubility curve

Is the graph of solubility versus temperature.

Graph curve

Action of Heat On Salts

Action of heat on carbonates.

Potassium carbonate and sodium carbonate does not decompose on heating.

2CO3 No effect

a2CO3 No effects

Other carbonates when heated produce oxide metals and carbon dioxide

gCO3( s) MgO(s)+ CO2(g)

aCO3(s) CaO(S) + CO2(g)

Action of heat on nitrates.

The effects of heat on nitrates depend on the arrangement of o in their reactivity series (electro chemical series) of metal.

Is the arrangement of metal in order of their power to reacts

table 3.6

K Na	Decompose to form nitrite and oxygen 2NaNO3 2NaNo2 + O2, 2KNO3 +2KNO2 +O2
Ca Mg Al Zn Fe Pb Cu	Decompose to oxide of metal, nitrogen dioxide and oxygen 2Ca (NO3)2 2CuO+ 4NO2 +O2 2CU(NO3)2 2CUO + 4NO2
Ag Hg Au	Decompose to nitrogen dioxide, metal and oxygen 2Ag No3(s) 2Ag + 2NO2 + O2 Does not exist

Ammonium nitrate decompose on heating to form dinitrogen oxide and water.

H4NO3(S) N2O +2H2O

Action of heat on sulphates

Sulphates are generally more stable to heat than nitrates even sulpahtes of metals low their reactivity series must be strongly heated to decompose

Sulphates of alkalis metals and alkali earth metals do ot decompose when heated.

Few sulphates decompose on heating

Ammonium sulphates decompose on heating yo give sulphuric acid and ammoniu gas.

During the sulphate first decomposes to give ammonium, hydrogen sulphates and ammonia gas.

(

NH4)2 SO4(s) NH4HSO4(s) + NH3

n cooling ammonia gas and sulphuric acid are formed

(

NH4) HSO4 NH3(g)+ H2SO4(l)

Action of heat on chloride

Generally all to form hydrogen chloride and ammonia gas.

Ammonium salt decompose on heating. Ammonium a chlorides sublimes On heating

H4Cl(s) HCl(g) + NH3(g)

Uses of Salts.

Control of soil PH.

Soil that too acidic or two alkaline are not suitable for crop production. Calcium oxide (lime) is usally added to acid soils to neutralize the acidity. When the soils alkaline calcium sulpahtes (gypsum) is added.

Uses of salts as antacids

Anticids are substances that are in form of hydroxides or salts. They are used to relieve hurt burn and acids build up in the stomach, some of common antacids are magnesium sulphates and sodium hydrogen carbonate

As inorganic fertilizers are salts. Examples of these are ammonium sulphates, ammonium nitrates and calcium phosphates.

Alleviation of health disorders.

Various salts are used to alleviate problem associated different conditions. These salts include sodium salts, (chloride, phosphate, sulphate) and calcium salts among others

Other uses.

Sodium chloride is used for seasoning and preserving food

Sodium carbonates is to soften hard water and making glass

Ammonium chloride is used as an electrolyte in the manufacture of dry battery.

Ammonium salts are used in the manufacture of plastic a synthetics fibers, dyes and explosives, and explosives and pharmaceuticals. Ammonium salts

Copper (II) sulphate issued as a fungicide.

Silver bromine is used in making photographic films.

Calcium chloride is used a drying agent and in freezing mixtures. Because has attended of absorbing water and it does not become a solution.

Terms Used

Hygroscopic

Salt which absorb moister and does not become solution the tendency is called hygroscope

Deliquescent

Salt which absorb moister and become a solution the tendency is called deliquescence.

Fluorescent

Are salts which does not absorb moister. The tendency is called florescence.

Acids Bases and Salts Summary

An acid is a substance that dissociate in water to produce H+ as the only ions

A base is a substance that ions in water to form hydroxide ions.

A strong acid ionizes completely in water to produce hydrogen ions. A weak acid ionizes partially in water

A strong base ionize completely in water while a weak base increases partially.

Water is a polar solvent so it can dissolve other polar substance. HCl can ionize in water but not in monthly benzene. HCl in monthly benzene does not conduct electricity because monthly benzene isnon polar.

A salt is an ionic compound and formed when a cation derived from a base combines with an anion derived from an acid

A reaction reaction resulting in formation of solids is called precipitation reaction

Solubility is the amount of solute that dissolve in of water a

To saturated it.

Water which does not form latter is said to be hard..

Water which contains calcium and magnesium hydrogen carbonate is saiud to have temporary hardness.

Water which contains calcium carbonate and magnesium carbonate is said ton be hard water.

Hardness can be removed using several ways

Substance which posses sour taste are said to be acid for example, the vinegar, sour milk they turn blue litmus paper red.

Substances which neutralize acids to form salts and water only are known as bases.

Acids react with various substances to give specific products for example they react with metal metals above hydrogen in the reactivity series to give salts and hydrogen gas is liberated.

The reaction of an acid and abuse to indicate whether a substance is an acidic or alkaline.

A salt is a compound which is formed when the hydrogen ions of an acid is replaced by a metal ion or ammonia ion.

END OF TOPIC QUESTIONS

Which of the following statements about solubility of salts is wrong?

Salt of K+, Na+, NH4+, NO-3 and HCO-3 are all soluble

Salts of Cl-, SO42-, CO2-3 and OH are all insoluble

Salts of Cl- are all soluble except those of Ag+, Hg+ and Pb2+

Salts of CO2-3 are all insoluble except those of Na+, K+ and NH4+

Which one of the following chemical substances shown by formulae form weak electrolyte?

NaCl

H2CO3

H2SO4

MgCl2

Which action should be taken immediately after concentrated sulphuric acid spilled on the skin?

It should be rinsed with large quantities of running water

It should be neutralized with solid CaCO3

It should be neutralized with concentrated NaOH.

The affected area should be wrapped tightly and shown to a medical health provider.

It should be neutralized with concentrated KOH

The ionic equation of the reaction between hydrochloric Acid and Sodium hydroxide is;

a+ + Cl- NaCl

a+ + OH- NaOH-

3O+ + Cl HCl + H2O

3O + Cl HCl + H2O

H+ + OH- H2O

A hygroscopic substance;

Evaporates to the air

Loses water to the air

Add water to a compound

Absorbs moisture from the air

Removes elements of water from a compound

An efflorescent substance;

Evaporates to the air

Loses water to the air

Absorbs moisture from the air

Removes elements of water from a compound

Adds water to a compound

The ionic equation when aqueous ammonium chloride reacts with sodium hydroxide solution is represented as;

NH4+(aq) + 2Cl-(aq) 2NH3(g) + +H2(g) + Cl2(g)

H+4(aq) + OH-(aq) NH3(g) + H2O(l)

+(aq) + OH(aq) H2O(l)

NH4(aq) + 2Cl(aq) 2NH3(g) + 2HCl(g)

H+4(aq) + 2Cl- (aq) 2NH3(g) + 2HCl(g)

The following are volatile acids except;

Ethanoic aid

Nitric acid (concentrated)

Carbonic acid

Hydrochloric acid

Only metal which does not react with dilute hydrochloric acid is;

Magnesium

Aluminium

Zinc

Sodium

Copper

An oxide of metal that react with both acid and base is known as;

Mixed oxide

Neutral oxide

Amphotesic oxide

Basic oxide

Peroxide

Which of the following statements is true about acids?

All acids ionize completely to give hydrogen ions.

All acids react with bases to form salt as of the products.

Acids change a blue litmus paper red.

Acids change a red litmus paper blue.

The following are true about bases except ……………

Most bases have better test

Bases are slippery when rubbed between fingers.

Bases neutralize acids to form salt and water only.

Weak base is the one that dissociates completely in solution to give hydroxide ions.

Methyl orange (M.O) changes colourless acidic solution to;

Orange B. Red C. Blue D. Pink

Which of the following salt is not soluble in water;

Sodium chloride

Potassium chloride

Lead chloride

Ammonium chloride

Most salt have comparatively high melting points because they have;

Crystalline structure

Low pressures

High specific heats

Strong specific heats

A hens egg shell contains …………………….

Calcium sulphate

Calcium carbonate

Calcium chloride

Calcium nitrate

Which one of the following chemical compound is used in car batteries?

Sodium chloride

Sodium hydroxide

Sulphuric acid

Ammonium sulphate

Which is the common physical property in the following chemical compounds? H2SO4, CaCl2.6H2O, KNO3, NaCl, CuSO4, (NH4)3 PO4, NaCO3. 10HO2 They are:

Soluble in water

Hygroscopic

Deliquescent

Efflorescent

Find solutions A – E have the following pH values: A. ph1 B. PH6 C. PH7 D. PH9 E. PH12

Which of the solutions:

Is exactly neutral

Is strong acidic

Is weakly alkaline

Could be a solution of carbon dioxide in water

Will give universal indicator colour (i) Green (ii) Red

Will contain a large concentration of OH ions

Could be a solution of ammonia in water?

Hard water which can be softened by boiling method contains dissolved:-

Calcium Hydrogen sulphate

Magnesium chloride

Magnesium carbonate

Calcium sulphate.

Which of the following compounds can be used to remove hardness of water?

Calcium chloride

Sodium carbonate

Magnesium hydroxide

Potassium hydroxide

Zinc chloride.

PART II

Hydrochloric acid ionizes in water as follows; HCl(aq) + aq → H+(aq)+ Cl(aq)

Define the term acid

HCl is said to be a strong acid , explain

Ammonia gas dissolve in water as follow

NH3(g) + H2O(l) → NH4(aq)+ +OH−

Define base

Ammonia is a weak base explain

Write equation for reaction between hydrochloric acid and ammonia solution

120g of saturated solution at 650C Contains 50g of potassium chloride

Define solubility

Calculate the mass of water used to make solution

Calculate solubility of potassium chloride in water

Complete the following equations

H2SO4 (aq) + Zn(s) →?

H2SO4(aq + Zn(s) →?

Use the information below on solubility and answer the questions that follows

Solubility	Solubility at different temperature	Solubility at different temperature
	At 800C	At 30oC
CuSO4	48	38
Pb(NO3)2	88	89

Which of the two salts will crystallize first when temperature drops from

800C to 300C.

calculate the mass that is formed

which salt is unsaturated at 300C

(a) Differentiate between the following terms;

Deliquescent against hygroscopic salts

Hydrated salt against anhydrous salt.

Water of crystallization against efflorescent substance.

Base against alkali.

(b) Explain briefly all categories of salts giving one example for each category.

(a) Distinguish between;

Acidic salt and basic salt

Hydrated and anhydrous salt

(b) Vinegar, lemon juice and yoghurt all have sour taste. Give the other possible properties would you expect them to have? Give 3 points.

(a) Complete and balance the following equations;

Fe(s) + CUSO4(aq)

Ca(HCO3)2(aq)

HNO3(aq) + NaOH(aq)

CaCl2 . 6H2O(s)

(NH4)2 SO4(s)

(b) A sample of water forms scum with soap. When the water was boiled and then cooled it still formed

scum.

Does this water have temporary hardness or permanent hardness?

Suggest one treatment that can be used to soften the water.

(a) Define the following words:-

Deliquescence

Efflorescence

(b) Give at least three (3) uses of salts

(c ) When zinc granules and dilute sulphuric acid are reacted together, gas M is produced. The gas produced is collected by downward displacement of water. Use this information to answer the questions below:-

Name the gas M

How is gas M tested?

Why is a the gas collected by down ward displacement of water?

Mention at least four (4) apparati used in laboratory preparation of gas M

(a) Define the term “basicity of an acid”

(b) Give any three (3) examples of monobasic acids

(c ) (i) write down the products formed when and acid reacts with an active metal.

(ii)Describe a test for identifying the gaseous product in c (i) above.

The basicity of Sulphuric acid is in 2 using equations explain what ismearnt by basicity.

Define hardness of water

(a).What ions cause I. temporary hardness of water ii. Permanent hardness of water

(b)Suggest the best methods that can be used to remove the

i. temporary hardness of water

ii. permanent hardness of water

©. The table below shows quality of soap needed for

Washing using two types of water

water	Time taken to clean	Soap used in (g)
x	8minites	50g
y	4 minutes	15g

Which water is hard .explain

Why do you think it took more time to clean using x than y

One way of removing water hardness is ion exchange explain how it work.

Felista dissolve HCl gas in methyly benzene and in water.

Explain what will happen when litmus paper is added into the two solutions.

17. Define (i) Temporary hardness of water and (ii) Permanent hardness of water

(b) Write balance equations of reactions used to solve hardness of water problems by (i) Heating when magnesium hydrogen carbonate is present in the water (ii) Precipitation using sodium carbonate to remove calcium ions due to the presence of calcium chloride in water

18.(a) Define (i) a deliquescent substance (ii) hydroscopic substance

(b) Which of the following salts are:

(i)Drying agents

(ii) Salt hydrates

(iii) efflorescent salts

Salts: CaCl2, Na2CO3. 10H2O, CuO, CaO

Acids can react with, carbonates, hydrogen carbonates, oxide and hydroxide. Presence of carbon dioxide can be proved using lime water. Support the above statements by using balanced chemical equations.

Distinguish between the following terms

Hydrate salts and anhydrate salts

Efflorescence and deliquescence

Base and alkali

Acid and base

Acid salts and normal salts

(a) Differentiate between;

A base and an alkali

Atom and Isotopes

(b) An organic compound P consists of 52.2% of carbon, 13% of hydrogen and 34.8% of oxygen. The vapour density of P is 23. Calculate the molecular formula of the compound P.

(a) Outline one use for each of the following;

NaOH

KOH

Co(OH)2

NH4OH

Mg(OH)2

(b) State whether each of the following is a strong Acid or weak acid;

i. Tarlartaric acid

ii. Ethanoic acid

iii. Hydrochloric acid

iv. Phosphoric acid

v. Methanoic acid

The word equation below shows the reaction of Zinc with dilute Sulphuric (VI) acid.

Sulphuric (VI) acid + Zinc → Zinc sulphate + Gas X. Identify gas X.

List the following substances either as acid, basic or neutral: wood ash solution, sour milk, lemon juice, toothpaste, vinegar, distilled water, common salt solution.

Give the products of the reaction between a dilute acid with:

Metal hydroxide.

Metal oxide.

Carbonate.

A bee sting is said to be acidic, how can it be treated?

Differentiate an acid from base in terms of PH.

Considering the PH scale, give an appropriate pH Values for the following solutions: lime water, lemon juice, sulphuric (VI) Acid, potassium hydroxide.

Given lead (II) oxide, how would you prepare lead (II) Chloride?

Describe how you would extract an acid base indicator from flower petals.

What would be the effect of adding lime to soil sample with pH value of 4.0?

List down any 3 uses of salts in everyday life.

For each of the following write equations for the reactions that would take place.

Action of heat on Sodium nitrate, Lead (II) nitrate and mercury (II) Nitrate.

Action of heat on sodium carbonate and copper (II) Carbonate.

Write the general equation for the reaction of an acid and a base.

Name the physical process used to obtain salt crystals from a salt solution.

Sulphuric acid is a dibasic acid and forms two series of salts with sodium hydroxide. Give an example of a salt for each of the two series.

The Ph value of 0.001M Sulphuric acid (H4 SO4) Solution is about the same as that of 0.1M ethanoic acid (CH3COOH) solution. Explain this observation.

TOPICAL QUESTIONS ON ACIDS, BASES AND SALTS

FORM THREE CHEMISTRY.

SECTION A. 20 MARKS.

1. MULTIPLE CHOICE QUESTIONS

Sea water contains various salts. Which salt is present in the largest proportion?

Magnesium sulphate

Sodium chloride

Calcium sulphate

Magnesium chloride

Which pair among the following comprises of good drying agents?

Calcium carbonate and lead (II) chloride

Anhydrous calcium chloride and concentrated sulphuric acid

Calcium sulphate and calcium oxide

Concentrated nitric acid and concentrated hydrochloric acid

A weak acid is the best describe as ……………………..

An acid that does not ionize completely.

A dilute acid.

An acid that does not react with any substance.

An acid that is harmless.

When dilute solutions of calcium chloride and sodium carbonate are mixed;

A white precipitate of sodium chloride is formed

A white precipitates of calcium carbonate is formed

A colourless solution of calcium carbonate and sodium chloride are formed

A mixture of precipitates of sodium chloride and calcium carbonate are formed.

Insoluble salts like lead (II) chloride, generally can be obtained in the laboratory by;

Crystallization

Precipitation

Decomposition

Evaporation of its concentrated solution

A hygroscopic substance;

Evaporates to the air

Loses water to the air

Add water to a compound

Absorbs moisture from the air

Removes elements of water from a compound

An efflorescent substance;

Evaporates to the air

Loses water to the air

Absorbs moisture from the air

Removes elements of water from a compound

Adds water to a compound

Which of the following is a weak acid?

Sulphuric acid

Hydrochloric acid

Ethanoic acid

Nitric acid

A solution of a PH of 1 is said to be;

Slightly acidic

Slightly basic

Strong acid

Neutral

Which of the following is not a product when silver nitrate is heated?

Silver oxide

Silver metal

Nitrogen dioxide gas

Oxygen gas

2. Matching items questions.

LIST A

LIST B

A salt that absorbs water to form solution

A salt that loses water into the atmosphere

A salt that absorbs water from atmosphere but do not form solution

Number of replaceable hydrogen ions

An acid whose basicity is three

A salt that is used in softening water

A metal nitrate which when heated decomposes to form metal nitrite and oxygen

A gas with a brown colour

An oxide which is orange when hot and yellow when cold

Reaction between an acid and a base to form salt and water only

Nitrogen I Oxide

Nitrogen dioxide

Lead oxide

Zinc oxide

Basicity of an acid

Delinquent

Hygroscopic substance

Efflorescent

Potassium nitrate silver nitrate

Hydrating salt

Drying agent

Calcium sulphate

Sodium carbonate

Neutralization reaction

Sulphuric acid

Phosphoric acid

Displacement reaction

How can one prepare a flower extract indicator and use it?

Distinguish between the following terms

Hydrate salts and unhydrate salts

Efflorescence and deliquescence

Base and alkali

Acid and base

Acid salts and normal salts

Fill the following tables

Table one.

Test	Observation
1. Blue and red litmus paper dipped into sodium hydroxide
2. A piece of magnesium metal dropped into diluted hydrochloric acid
3. One spatula of sodium carbonate added into dilute sulphuric acid
4. 2cm3 of AgN03 solution added to the solution of mgcl2

Table Two.

Indicator	Colour change in acidic solution	Colour change in alkaline solution
1. Phenolphalein
2. Methyl orange
3. Litmus paper

(a) 416g of anhydrous barium chloride where obtained when 488g of hydrated salt were heated.

Calculate the value of n is the formula BaCl2.nH2O (Ba = 137)

(b) Give the meaning of the following;

(i) An acid

(ii) Molar solution of an acid

(iii) Give three characteristics of acid.

Differentiate salts behave differently when heated. Complete the following equations that show the actions of heat on certain salts;

bCO3(s)

Ca (NO3)2(s)

(

NH4)2 CO3(s) 2NH3(g)+

H4Cl(s)

(a) Distinguish between strong acid and weak acid

b) Explain three applications of neutralization

From above information answer the following questions;

(i) What piece of apparatus should be used to measure out accurately 25cm3 of sodium hydroxide solution?

(ii) What color was the solution in the flask at the start of the titration?

(iii) What color did it turn when the alkali had been neutralized?

(i) Was the acid more concentrated or less concentrated than the alkali?

(ii) Name the salt formed in the neutralization.

(iii) Write an equation for the reaction.

(iv) Is the salt, normal or acidic salt? Give reasons for your answer.

(i) Utilizing the given information of question 9 (a) and (b) above describe how you can obtain pure crystals of the salt.

(ii) Is it a soluble salt or insoluble salt? Give a reason for your answer.

A sample of water forms scum with soap. When the water was boiled and then cooled it still formed scum.

Does this water have temporary hardness or permanent hardness?

Suggest one treatment that can be used to soften the water.

(a) Define the term indicator.

(b) Write only two uses of pH scale.

(a) Give meaning of the following;

Basicity of an acid

The pH scale

(b) Name substances which when dissolved in water causes;

i. Temporary hardness of water

ii. Permanent hardness of water

i. Temporary hardness of water can be removed by boiling.

ii. Permanent hardness of water can be removed by chemical means.

(a) Differentiate between basic salt and acidic salt.

(b) Categorize the following salts;

i. PbSO4 ii. MgHSO4 iii.Zn(OH)Cl iv. MgHCO3

v. NH4HSO4 vi. Ba(NO3)2

Salts	Category

(a) Define the term “neutralization reaction” (give one example)

(b) Write down the names and formulae of three common acids in the laboratory.

(c ) What is an indicator? Give four (4) examples of acid- base indicators.

(d) Write down the products formed when each of following pairs of compounds react:

(i) Acid and metal

(ii) Acid and metal carbonate

(a) Name of element found in all acids

(b) Describe two tests you would carry out on a liquid to decide whether it is an acid or not in each case describe the test and result you would expect.

SAMPLE NECTA QUESTIONS

2006.

b) Define the following terms and give one (1) example in each case:

Weak acid

Acidic salt

c) Write ionic net equation for the following reactions:

(1) Barium chloride when reacting with sodium sulphate.

(ii) Sodium hydroxide when neutralizing hydrochloric acid.

Calcium metal when reacting with dilute hydrochloric acid.

2009.

QN.3 Why doesnt water have any effect on litmus paper?

(i) What would happen to a well stoppered bottle full of water left in a deep freezer over night? Why does this happen?

(ii) Why isnt iron usually recommended in the construction of steam pipes and boilers? Explain.

(i) Name two ions which cause temporary hardness of water and two ions which cause permanent hardness.

(ii) Give equations for two ways to remove temporary hardness of water and one way to remove permanent hardness.

2011.

QNS.5 (a) Giving four reasons, explain why people who use hard water can experience higher costs than people who use soft water.

(b) Suggest one method for the separation of each of the following: (1) Iodine and sand

Green solution from leaves

Alcohol and water

Iron fillings and powdered calcium carbonate

TOPIC : 4 FUELS AND ENERGY

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

ENERGY AND FUELS

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

MULTIPLE CHOICE QUESTIONS.

Which one of the following energy transformation can produce H.E.P?

Electrical energy changes to heat energy

Electrical energy changes to mechanical energy

Mechanical energy changes to Electrical energy

Mechanical energy changes to chemical energy

One of the following is not correct about coke being a better fuel than coal as it;

Does not produce carbon dioxide gas

Does not produce poisonous gas

Has a higher heat content [ ]

Is clean and smokeless

___________is the characteristic of good fuels:-

Firewood

Easily available ( )

Melting

Ability to burn houses into ashes.

What is kindling temperature

A kind temperature

Temperature out of a burning material

The highest temperature obtained from a burning substance

The lowest temperature at which a combustible material can catch fire.

Which of the following is a renewable source of energy?

Coal

Petroleum

Biogas

Natural gas

Which of the following is not a primary source of energy?

Wood

Kerosene

Crude oil

Natural

Which of the following is not a characteristic of a good fuel?

Readily available

Should have low energy value

Should be affordable

Easy to transport

In production of hydroelectric power, which of the following is the best energy transformation?

Electric energy to heat energy

Electrical energy to mechanical energy

Mechanical energy to electrical energy

Chemical energy to electrical energy

Which of the following has a chemical energy?

Natural gas

Biomass

Coal

Solar energy

Which of the following is not an advantage of biogas?

It is cheaper source of energy

Pollutes the environment

It is renewable source of energy

Creates employment among the youth

Matching items questions

LIST A

LIST B

Combustible material that gives large amount of heat.

Fuel that does not pollute the environment

Fuels that occur in nature

Fuel that is derived from primary fuel

Temperature at which fuel must be heated before it starts burning

Highest temperature that can be reached by a burning fuel

The rate at which a fuel burns

Capacity to do work

Fuel formed from remains of plants and animal materials.

Renewable source of energy from waste

Pyrogalic burning effect

Kindling point

Clean fuel

Polluting fuel

Primary fuel

Natural fuel

Secondary fuel

Velocity of combustion

Energy

Power

Biomass

Biogas

Wind

Non-combustible materials

Electricity can be used to perform work also for heating purpose. Explain how it can happen and use four (4) practical examples in each case.

(a) What is fuel? ……………………………………………………………………………………………………………………………………………………………………………………………………

(b) Mention three categories of fuel and give two examples for each category.

(i) ………………………………………..

examples …………………………………. and ………………………………….

(ii) …………………………………………

examples ………………………………… and …………………………………

(iii) ……………………………………………

examples …………………………………. and ………………………………….

source of fuel. Give two points of advice to the society on how to use less

charcoal and firewood efficiently.

(a) Write P for primary and S for secondary in the following fuels.

Wood……………. Petrol…………… Coal gas……………………….

(b) Which gaseous fuels are a result of the following processes?

(i) Destructive distillation of coal…………………………………….

(ii) Destructive distillation of wood…………………………..........

(iii)Air reacts with red hot coke at 10000C……………………..

(a) Write the chemical symbol for each of the following elements.

Argon_________(ii)Lead________(iii)Silver__________(iv)Manganese____ (v) Barium______________(vi)Xenon_____________.

(b) Write the chemical formulae of the following compounds:-

Potassium sulphate

Copper (II) carbonate

Aluminium oxide

Magnesium hydroxide

Explain the major ways of obtaining artificial fuels from natural fuels in all the three states of matter.

By the help of a well labeled diagram explain how one can conduct an experiment on destructive distillation of coal.

(a) What is fuel

(b)Give two example of each of following categories of fuels

i) Solid fuel

Liquid fuel

iii) Gaseous fuel

( c) State any two qualities of good fuel

(a) Define the following terms

Green house effect

Global warming

Acidic rains

Pollution

(b)Give any four examples of green house gases

(a) Define the following terms:

Fuel ……………………………………………………………………………………………………………………………………………………

Calorific value of a fuel ……………………………………………………………………………………………………………………………………………………

Energy value of a fuel ……………………………………………………………………………………………………………………………………………………

(b) Give two examples of each of the following:

(i) Solid fuel ………………………………….., ……………………………..

(ii) Liquid fuel…………………………………..,……………………………...

(iii) Gaseous fuel ……………………………….., …………………………….

(i) ………………………………………………………………………………

(ii) ………………………………………………………………………………

(iii) ……………………………………………………………………………..

(iv) ……………………………………………………………………………..

School Base-OnlinePage 5

TOPIC : 5 MOLE CONCEPT

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

MOLE CONCEPT

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

i) If 0.9g of calcium metal is burnt inair, the mass of powder formed is:

1.14g

1.18g

1.12g

1.26g

ii) If 1g of hydrogen is exploded in air, the mass of water formed is:

1.8g

18g

What mass of pure sulphuric acid is found in 400cm3 of its 0.1M?

2.45gm B. 9.80gm C. 3.92gm D. 4.90gm

The volume of 18M concentrated sulphuric acid that must be diluted with distilled water to prepare 10 litres of 0.125M sulphuric acid is;

69.44cm3 B. 22500cm3 C. 225cm3 D. 4440cm3

If two jars labelled W and Z contain 22.4cm3 of oxygen gas and 22.4dm3 of nitrogen gas at STP respectively, then it is true that;

There were 6.02 x 1023 oxygen molecules in jar W and 6.02 x 1023 nitrogen molecules in jar Z.

6.02 x 1023 oxygen atoms were in jar W and 6.02 x 1023 atoms of nitrogen in jar Z.

There were 12.4 x 1023 molecules of oxygen and nitrogen in the gas jars W and Z.

6.02 x 1023 molecules of oxygen and nitrogen were in the two jars W and Z.

Which of the following statement is true about one mole of CO2;

It is contains Avogadro’s number of molecules.

It contains 6.023 x 1023 molecules at room temperature

It contains 44g of carbon and 16g of O2

It contains 12g of carbon and 16g of O2

The loss in mass when 100g of marble is heated to a constant mass is;

200g B. 88g C. 44g D. 22.4g E. 22400g

The number of atoms present in 0.025 moles of iron is;

6.02 x 1023 atoms B. 56 atoms C. 5.6 x 1023 atoms D. 22.4 x1023 atoms E. 1.51 x 1023 atoms

How many atoms of zinc are there in 165g of zinc?

1.505 x 1025

1.0997 x 1022

2.2586 x 1023

9.63 x 1025

SECTION B

3.5 g of a hydrated salt, MSO4.XH2O, was heated to a constant mass of 3.21g of the anhydrous salt. Calculate the value of X. (M= 63.5, S= 32, O = 16, H= 1)

With the help of a chemical equation, calculate the volume of carbon (IV) oxide produced when 20g of calcium carbonate is heated at standard temperature and pressure (STP). (Ca= 40, C= 12, O =16,)

What would be the volume of the gas produced at 200C and 750mmHg.

Molar gas volume at STP is 22.4 litres)

25.0cm3 of sodium hydroxide solution containing 4.0g/dm3 were required for complete neutralization of 0.18g of a dibasic acid. Calculate the relative molecular mass (RMM) of the acid. (Na=23, H= 1, O=16).

(a) How many molar volumes of 132.0g of CO2 are there at STP?

(b) Determine the number of molecules in 0.25 moles of lead (II) nitrate.

(a) What mass in grams of hydrated sodium carbonate (Na2CO3 . 10H2O) in 65cm3 of 0.2M solution?

(b) What volume of carbon dioxide would be evolved at STP when 6.2g of copper (II) carbonate is reacted with sulphuric acid?

When excess lead (II) nitrate solution was added to a solution containing sodium chloride, a white precipitate was formed and weighed 5.56g. determine the amount of sodium chloride in the solution. (Bb = 207, Cl = 35.5, Na = 23)

32.5 of IRON (III) chloride was reduced by hydrogen to iron metal. Calculate the volume of hydrogen used to in the reaction at room temperature and pressure. The equation of the reaction is shown below.

2FeCl3(s) + 3H2 (g) → 2Fe(s) +6HCl (g)

(Fe = 56, Cl = 35.5, molar gas volume = 24 liters at r.t.p).

7. (a). How many moles of calcium are present in 20g of the metal?

b. determine the grams of magnesium that contains 1.8 × 1024 atoms

c. how many ions are there in 4.6 of Na+ ions? (Ca ═ 40, Mg ═ 24, Na ═ 23)

8. Find the mass of 3.0g of calcium carbonate.

150g of Calcium carbonate were reacted with 0.2 moles of hydrochloric acid.

Write equation of reaction

Calculate amount of calcium carbonate that will remain unreached.

0.7 g of gas × occupies 560 cm3 at s.t.p. determine its relative molecular mass.

Molar gas vol. = 22.4L

A certain carbonate aCO3 reacts with dilute hydrochloric acid according to the equation given below.

ACO3(s) + 2HCl (aq) → Al2aq + CO (g) + H2O (l)

If 1g of the carbonate reacts completely with 20cm3 of the IM HCl. Calculate the relative atomic mass of G. (C = 12.0, O = 16.0)

When a hydrocarbon was completely bunted in oxygen, 4.2g of carbon dioxide has1.71g of water were formed. Determine the empirical formula of late hydrocarbon. ( H= 1.0, C = 12.0 , O = 16.0 )

9. Suppose you have 3.92 g of hydrated ammonium iron (II) sulphate, how many moles of the compound do you have?

(H ═ 1, N ═ 14, Fe ═ 56, S ═ 32, O ═ 16)

10. How many grams are there in 3 moles of water?

11. Determine the number of molecules in:

1.3 dm3 of hydrogen at S.T.P

1.2 dm3 of sulphur (IV) Oxide at R.T.P

(L ═ 6.0 × 1023, Molar gas volume at S.T.P ═ 22.4 dm3 )

School Base-Online Page 4

TOPIC : 6 ELECTROLYSIS

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

ELECTROLYSIS

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

SECTION A. 20 MARKS.

1.MULTIPLE CHOICE QUESTIONS

During electrolysis the molten Aluminium oxide, 3 Faradays were needed to deposit one mole of aluminium. The number electrons of aluminium will be :

6.02 x 1023

1.806 x 1023

18.06 x 10.23

180.6 x 1023

1806 x 1023

If a steady current of 2 amperes were passed through an aqueous solution of iron deposited are the cathode will be:-

54g

56g

0.54g

28g

0.52g

The use of electricity to decompose molten sodium chloride into its components elements is an example of:-

Electroplating

Galvanization

Hydrolysis

Precipitation

Electrolysis

What is the meaning of electrolysis?

Decomposition of a chemical compound by electricity

Breaking up by using electric power

Oxidation reduction process

Passing electricity through a conductor

Which is the correct property of an electrolyte

It contains free charge carriers

It does not have free ions

It can not be decomposed by an electric current

It contains mainly undissociated molecules

Which of the following is not an electrolyte

Aqueous sulphuric acid

Solid common salt

Molten calcium chloride

Alcohol in water

When dilute sodium chloride solution is electrolysed using carbon electrodes which substance is collected at the cathode

Hydrogen

Sodium

Chloride

Oxygen

What are cations?

Cathode going ions

Anode moving ions

Negatively charged ions

Unionized particles

Which particles will be anions in the electrolysis of aqueous copper sulphate using platinum electrodes?

SO42- and OH

Cu 2+ and H+

Cu 2+ and SO4 2-

Platinum

What is the name given to quantity of substance deposited by one coulomb of electricity on the electrode?

Electrochemical equivalent

Chemical equivalent

1 mole of the substance

1 G.M.V at s.t.p.

Which one of the following is the application of electrolysis

Electroplating

Isolation of elements

Purification of metals

All of the above

What is the name of the vessel in which electrolysis process is conducted

Voltameter

Voltmeter

Avometer

Beaker

Match the items in list A with those from list B by writing the letter of the correct response from List B besides the item number in List A.

LIST A

LIST B

Anode

Cathode

Electrolyte

Electrolysis

Electrodes

Non- electrolyte

Strong electrolyte

Weak electrolyte

Chemical equivalent

Electrolytic reduction

Decomposition of substance in molten or in solution by passage of electric current

A reduction process by the use of carbon monoxide or hydrogen

Relative concentration of ions in solution.

Two poles of carbon or metal dipped in solution allow electrons enter or leave electrolyte

Coating metal with a thin coat of another metal by electrolysis

Positively charged electrode which electrons enter the external circuit.

Decomposition of chemical

Substance which will conduct electricity.

Hydrogen gas liberated due to the passage of electric current

Negative electrode

Extraction of most reactive metals

Oxygen gas in collected

Splits up during electrolysis

Faraday constant divide by its valency

Atomic mass divide by its valency.

Completely dissociate during electrolysis.

Partially dissociate during electrolysis.

Copper metal is deposited

Relative position in electrochemical series.

Substance that do not conduct electricity.

(a) Define

I) Electrochemical equivalent

Chemical equivalent

(b) An electric current was passed in series through solutions of calcium chlorides and copper (ii) sulphate. Carbon electrodes were used in both electrolytes. If 2.5 litres of chlorine gas measured at S.T.P. were produced, what volume of oxygen gas would also be produced? What mass of copper was produced?

(a) Explain the meaning of the following terms;

Electrolysis

Electroplating

(b) State faraday laws of electrolysis.

(c) (i) Write chemical equation for the reaction that take place at each electrode during the electrolysis of dilute sulphuric acid using platinum electrode.

(ii) 19300 coulombs of electricity was passed through a solution of copper (II) nitrate.

Calculate the mass of copper deposited at the cathode.

(a) By the help of a well labelled circuit diagram explain the electrolysis of a dilute

solution of sodium chloride using graphitic carbon electrodes.

(b) Is the above electrolysis having special name? If yes explain.

dryness of the electrolyte? Explain your answer.

Calculate the mass of magnesium metal that will be produced during the electrolysis of molten magnesium chloride if a current of 1.93A is passed through for 16 minutes and 40 seconds.

(a) State faradays 1st law of electrolysis

(b) What mass of silver and what volume of oxygen (at s.t.p) would be liberated in electrolysis by 9650 coulombs of electricity?

7. I)(a) State the faraday law’s of electrolysis

(b) An element P has relative atomic mass of 88. When current of 0.5A was passed through fused chloride of P for 32 minutes and 10 seconds, 0.44g of P was deposited at cathode;

Calculate the number of faraday required to liberate 1mole of P.

i).Write the formula of P ions

ii).Write the formula of hydroxide of P

(a) Distinguish between chemical equivalent and electrochemical equivalent of a

substance.

(b) Describe electrolysis of copper (II) sulphate using graphite electrodes.

Note: Diagram is necessary.

I) (a) State Faraday laws of electrolysis.

(b) An element Z has a relative atomic mass of 88. When a current of 0.5 amperes was passed through fused chloride of Z for 32 minutes and 10 seconds 0.44 of Z were deposited at the cathode;

Calculate the number of Faradays needed to liberate one mole of Z.

Write the formula of the Z ions

Write the formula of hydroxide of Z.

(a) i. List down three (3) factors affecting the selection of ion discharge at the electrode.

ii. Define the term electrolyte.

(b) A bluish copper sulphate aqueous solution was electrolysed by using copper

electrodes.

Write ionic chemical equations for the reactions which occurred at the cathode and anode.

Explain what will happen to blue colour of copper sulphate solution as electrolysis continues.

The following apparatus was used in experiment to electroplate an iron knife with Silver;

The K.Ag(CN)2 contains Ag+, K+ and CN ions;

ii. Which electrode is the cathode?

Name the process taking place at the anode as either oxidation or reduction.

Represent the process at each electrode by the appropriate ionic equation.

What happen to the electrolyte?

(a) 0.02 moles of electrons were passed through a solution of sodium hydroxide using platinum electrodes.

Give the names of the gases evolved to each electrode.

Write ionic equations of the reaction taking place at the electrodes.

Calculate the number of moles of each gas produced and the volume which each gas would occupy at S.T.P.

(b) What mass of copper will be liberated during electrolysis of copper sulphate solution by a charge of one faraday?

(c) An element X has a relative atomic mass of 88. When a current of 0.5A was passed through the fused chloride of X for 32 minutes 10 seconds, 0.44g of X was deposited at the cathode.

Calculate the number of faradays needed to liberate 1 mole of X.

Write the formula for X ion.

Write the formula for hydroxide of X.

School Base-OnlinePage 6

TOPIC : 6 VOLUMETRIC ANALYSIS

CHAPTER 6

VOLUMETRIC ANALYSIS

KEY CONCEPTS AND TERMINOLOGIES.

Volumeric analysis- This is a chemical procedure for determining the concentration of a solution.

Standard solution- this s solution whose concentration is clearly known. During volumetric analysis, a known volume of a solution of unknown concentration is reacted with a known volume of a solution of known concentration.

Titration- this is the process by which a standard solution is delivered from a burrete so that the volume added is measured.

Indicator- this is a chemical substance that is used to indicate a change in colour when correct proportions has reacted.

End point- this is the point at which proportions would have reacted completely.

Analyte -is the reagent in the reaction vessel and whose concentration is to be determined.

Standardization- is the process used to determine the concentration of a second solution whose concentration is unknown.

Primary standard- is a pure substance with high degree of purity that is used to prepare secondary standard.

Secondary standard- this is a solution that is standardized using primary standards.

Dilution- this is a process by which the concentration of the solution is reduced by adding water.

Titration- this is a laboratory technique by which we determine the concentration of a solution, using a standard solution.

The basic principles of volumetric analysis are:

The one solution to be analyzed contains an unknown amount of chemicals.

The reagent of known concentration reacts with a chemical of unknown amount in the presence of an indicator which shows the end-point. This is the point at which the chemical reaction between the two reagents is complete.

The amount of unknown chemical in the measured volume of solution is calculated using the mole ration from reaction equation.

The amount of unknown chemical in the original sample is calculated by finding out the amount of unknown chemical in the measured volume.

The amount of substance present in a solution is given in terms of its mass in grams while the volume of a solution is usually given in litres or dm3. The concentration of a solution is given in moles per litre (mol/dm3) or mass per litre.

Hence:

Concentration is the amount of a solute dissolved in one litre of a solution

Thus: C =

where n = mole, v = volume in litres.

Concentration =

The reagents commonly used are acids and bases. The commonly used acids are:

Hydrochloric acid,

Sulphuric(VI) acid and

Nitric(V) acid while the commonly used bases are

Sodium hydroxide,

Potassium hydroxide and

Sodium carbonate.

Importance Of Volumetric Analysis.

Volumetric analysis is very important in determining the concentration of unknown solution, which has a lot of application in the industry.

Standard volumetric apparatus

Volumetric apparatus

These are apparatus used in carrying out volumetric analysis.

They include, a burette, a pipette, a retort stand, a white tile, a conical flask, a filter funnel, reagent bottles, an indicator, a wash bottle, a watch glass, beakers, a measuring cylinder and a volumetric flask.

The burette

It is used to measure accurate volume of the titrant (acid) during titration. Before use, the burette has to be rinsed with the solution to be measured. After use, the burette should be rinsed thoroughly with distilled water. Ensure that the burette does not leak before using it. The burette is graduated (calibrated) from 0cm3 to 50 cm3.

How to use a burette

Wash the burette with distilled water, and rinse it with the titrant solution.

Clamp the burette and use a filter funnel to pour the titrant solution into the burette while holding the funnel with one hand to allow air to escape.

When the burette is filled slightly above the 0cm3 graduation mark, remove the funnel.

Allow some solution to run out to ensure the air space below the tap is filled with the solution.

When titrating, the volume added is obtained by subtraction of the initial reading from the final reading.

The pipette

This is an apparatus used for measuring a fixed volume of a solution. It consists of a narrow glass tube into which small amounts of liquid are suctioned for transfer. It capacity is usually marked on its stem, with a certain mark.

The capacity of the pipette varies. We have pipette with capacities of 20cm3 and 25cm3. You can fill the pipette by sucking the solution with mouth or by using a pipette filler to help in the suction of a liquid into a pipette.

How to use a pipette

Wash the pipette with distilled water, then rinse it with the analyte.

Suck the analyte up the pipette until the solution is slightly above the etched mark. You can use your mouth or pipette filler.

The conical flask

Is a conical shaped glassware that is used to mix solutions during titrations. Its shape ensures that the solution does not spill during swirling.

Conical apparatus are available in various capacities such as the 250cm3 conical flask.

Molar Solutions

This is the number of moles of a substance dissolved in one litre of solution.

A molar solutioin is when one mole of a solute s in one litre of solvent. Written as M.solution

3 moles dissolved in one liter the concentration becomes 3 Molar.

The concentration of solution can also expressed in glucosen glucose of solution i.e conc = g/ litre.

E.g 40g of sodium hydroxide concentrated in litre ofsolution will have a concentration of 4g/litre. This is equivalent to 0.1M.

Example .

Given the molar mass if sulphuric aci is 98g. what mass of thin sold will be contained in 1M, 0.1m & 0.2m?

Solution

1mole→ 98g

Hence 1M = 98g

1Mole− 98g

= 98×0.1= 9.8g

0.2= 0.2× 98=

19.6g

The morality of solution is a term used for concentration

A M solution contains a moles of solute in litre

e.g 0.2M= will have 0.2 moles on litre

formula

concentration=

concentration =

moles of solute

we can also get the mass per litre as follows

mass per cm3= Molarity × Molar mass

Example

Calculate the solution containing 21.25 nitrate of sodium(NaNO3) in 2dm3 of solution.

Working

R.M.M of NaNO3 = 23+ 14 + 48

= 85.

Moles of NaNo3 21.25g =

= 0.25mol

Concentration =

= 0.125 M

Or if 21.25g are in 2 liter

10.625g → 1 Litre

Molarity =

= 0.125M

Example 2

Calculate the number of moles solute in 600cm3 of 0.5M sulphuric acid solution.

Answer

Moles of H2SO4 = Molarity × volume

= 0.5 × 0.6

= 0.3 mole

0.3 = 0.3 × 98

= 29.4g

Example 3

What mass of aluminium will react with 50cm3 of 0.4M HCL?

Solution

2AL + 6HCL → 2AlCl3 + 3H2

2 : 6 : 2 : 3

Mole of HCl

= 0.4 ×

= 0.4 × 0.05

= 0.002 Moles

Moles of Al =

= 0.006

Moles of Al = 0.006

Example 4

Calculate the morality of solution of sodium carbonate containing 1.06 g in 50cm3 solution.

Solution

Molar mass Na2Co3 = 2× 23 + 12 +16 × 3

= 106 g

Mass of Na2Co3 per liter

= 21.2 g

Molarity =

= 0.2 M

Example 5

Calculate the number of moles of alluminium in 20cm3 of 0.8M

Soluition:

Moles = Morality × volume]

= 0.0014

Standard solution.

This a solution whose concentration is known. It is used to establish the concentration of other solutions.

The concentration of a standard solution is expressed in moles per litre (mol/dm3). It is usually indicated by letter M for molarity.

In volumetric analysis, one of the solutions must have a known concentration. Volumetric analysis is used to determine the concentration of the second solution. This process is referred to as standardization. The second solution which, is being standardized is called the analyte.

Primary standards

Standard solutions are usually prepared from primary standards.

Primary standards should have the following characteristics:

A high degree of purity.

Stable, that is, they should not decompose with time.

Have no water of hydration, that is, they should not be hygroscopic or efflorescent.

Not volatile so that losses due to evaporation do not occur.

Be highly soluble.

Have a high molecular mass.

We can determine the amount of substance (number of moles) in a primary standard by simply weighing the substance. Examples of primary standards include:

Sodium carbonate (Na2CO3) for acid-base titrations, and

Potassium dichromate (K2Cr2O7) for redox titrations.

Secondary standards

Acids and bases such as dilute hydrochloric acid (HC1), dilute sulphuric acid (H2SO4), and sodium hydroxide (NaOH) are not primary standards, because they absorb water or vaporize during storage or preparation. Their exact amount cannot therefore be determined by direct measurement (weighing). However, their concentration in solutions can be determined by standardizing them with primary standards. The standardized solutions are referred to as secondary standards.

Preparing standard solutions

Standard solutions are prepared by dissolving a known amount of the primary standard in a known volume of a liquid. Pure water or distilled water is used for aqueous solutions.

If a particular primary standard is not available, the solution should be standardized with a primary standard of another substance. For example, sodium chloride (NaC1) is a primary standard which can be used to standardize a silver nitrate solution by titration.

The following is the procedure followed when preparing primary standards.

Calculate the number of moles of the substance needed to make up the standard solution. The required volume and concentration should be taken into consideration.

Work out the molar mass of the substance.

Using the results of the calculations in steps 1 and 2, calculate the mass of the substance needed to prepare the standard solution.

Carefully weigh the required mass of the substance. To weigh the required mass;

Weigh an empty watch glass. Add the mass of the watch glass to the mass of the required substance.

Weigh the watch glass together with the required mass of the substance.

Suppose 9.8 g of the substance required to prepare a substance and the mass of a watch glass is measured to be 75 g. The correct mass of the substance and the watch glass that should be measured on the weighing balance is 9.8 g + 75 g = 84.8 g.

Transfer the primary standard into a beaker. Rinse the watch glass using distilled water from a wash bottle. Ensure that the rinsing water goes into the beaker containing the substance being used to prepare the primary standard.

Stir the mixture in the beaker with a glass rod until all the solute is dissolved, then transfer the solution into a volumetric flask.

Rinse the beaker and the glass rod using distilled water from a wash bottle. Do this at least twice, and pour the rinsing water into the volumetric flask each time.

Add water to the solution to just below the graduated mark on the volumetric flask.

Top up the solution with water up to the graduation mark using a clean dropper. Ensure that the lowest part of the meniscus is exactly at the graduated mark.

Stopper the flask and invert it several times to ensure that the solution is homogenous.

The molar mass of sulphuric acid is 98 g/mol.

The number of moles of sulphuric acid in 9.8 Will be;

So, 0.1moles of sulpuric acid were dissolved in 1 liter.

Molar concentration of acid is 0.1M.

Preparation of standard solutions of common bases

Common bases such as sodium hydroxide are unlikely to have a high degree of purity. This makes it impossible to use them as primary standards.

To prepare standard solutions of the bases, they must first be standardized with acids that are primary standards.

Preparing standard solutions of common mineral acids

Standard solutions of mineral acids are prepared by dilution of the concentrated acids.

Concentrated acids from chemical shops are called stock solutions. The molar concentration of the stock solutions is usually not indicated on the reagent bottles. Only the density and percentage purity of the acids are indicated.

The volume of the flask is then adjusted to the etched mark, stoppered and thoroughly mixed by inverting the flask several times. The molarity of the dilute acid is then indicated on the flask.

This is only an approximate molar concentration of the acid. The acid is then standardized using a standard solution of a base for accurate determination of its molarity.

A bottle containing concentrated sulphuric acid could have the following information.

Relative molecular mass = 98

Density of acid = 1.84gcm3

Percentage purity = 98%

To prepare 1000 cm3 of 0.1M sulphuric acid solution from the sulphuric acid, follow the procedure below;

1 M sulpuric acid contains 98g in 1 liter or dm3

Thus, 0.1M sulphuric acid will contain, 98 x 0.1 mol = 9.8g in 1dm3

But density of acid is 1.84g which means 1.84g are in 1cm3

From percentage purity, 1cm3 of concentrated acid contain

Therefore, volume containing 9.8g of the acid is given

The following procedure is used in preparing the dilute solution.

Put 100 cm3 of distilled water in a beaker. Add about 5.4 cm3 of the concentrated sulphuric acid in a beaker.

Pour the solution into about 700 cm3 of cold distilled water in a 1000 cm3 volumetric flask.

Carefully add distilled water to the solution up to the 1000 cm3 mark.

Invert the flask several times to obtain a uniform solution.

The solution is approximately 0.1 M sulphuric acid. If the concentration of the original solution is C1 and that of the dilute solution is C2 the following relationship holds:

C1 V1 = C2V2

where V1 and V, are the volumes of the solutions of the respective concentrations.

This relationship is applied when preparing solutions of dilute solutions from concentrated solutions.

Task.

1.You are given the following information concerning dilute nitric acid;

RMM = 63 g

Density 1.42g/cm3

Percentage purity = 68%

Prepare a 500cm3 0.2M from the stock solution.

Use the dilution formular, C1V1=C2V2 to prepare a solution of concentrated hydrochloric acid of 0.4M from 100cm3 of 2M HCl.

Percentage purity of acids and alkalis

To find the percentage purity of a substance, calculate the amount (number of moles) of the pure chemical that would be needed to make the required solution, then, Multiply the number of moles of the pure substance by the molar mass of the substance to get the mass of the pure substance that is required.

That is, mass of pure substance = moles of pure substance x molar mass.

We can also find the mass of impure substance required to prepare a certain mass of pure substance by diving the mass of pure substance by percentage purity.

Sample questions

Potassium hydroxide has a purity of 85.9%, find the mass required to prepare 1 litre of 0.25M solution given that, mass of KOH is 56.11g/mol?

Find volume of HCl which is 37.1% pure and has a density of 1.2g/cm3 needed to prepare 1liter of 0.2M HCl, iven that mass of HCL is 36.5g/mol.

Titration

Titration is a laboratory technique by which we can determine the concentration of a solution using a standard solution. The two reagents should react chemically until an end point is reached.

For a chemical reaction to qualify for titrations, following are some conditions necessary,

The reaction must be fast, so that the titration can be performed in convenient time.

The reaction should go to completion, that is, it should not be a reversible reaction.

The reaction should be free from side reactions, meaning it can be represented by a single chemical equation.

The reaction should have a definite end point that can be determined accurately.

Common titration reactions include acid-base reactions, precipitation reactions and redox reactions.

Acid-base titrations

The most common titrations are acid-base titrations. Typically, the titrant (the solution whose concentration is known) is added from a burette to a known quantity of the analyte (the solution whose concentration is not known) until the reaction is complete. The completion of the reaction is indicated by the change in the colour of an indicator.

The indicator to be used depends on the concentration of the solutions in question. This is summarized in table below.

Titration reaction	example	Most suitable indicator
Weak base + weak acid Strong base + weak acid Strong base + strong acid Strong acid + weak base	Ethanoic acid + ammonia solution Sodium hydroxide + ethanoic acid Hydrochloric acid + sodium hydroxide. Hydrochloric acid + ammonia solution	No suitable indicator Phenopthalein and methyl orange Any indicator Phenopthalein or methyl orange

Table 5.1

Procedure for carrying out acid-base titrations.

The following are the steps followed whet performing acid-base titrations.

Clean the burette and the pipette before using them

Use a retort stand to hold the burette upright Fill the burette with the titrant, usually the acid.

Read and record the initial burette reading.

Pipette the alkali (base) into a conical flask.

Add one or two drops of the indicator to the alkali in the flask and note the colour of the solution.

Drain the acid slowly from the burette into the alkali, while carefully swirling the contents the flask, until a permanent colour change the indicator is observed.

When the end point is near, the colour of the indicator will start to change. At this point, add the acid in small drops. Note the point at which there is a permanent colour change in the reaction mixture.

Note the final reading of the burette.

Work out the volume of the acid used by subtracting the initial burette reading from the final burette reading. Record the volume of the acid used.

Repeat the procedure three more times to obtain consistent values which are within ± 0.2 cm3 of each other.

Using your results, work out concentration of the substance whose concentration is not known.

Standardization of hydrochloric acid using 0.098M sodium carbonate solution.

In order to standardize hydrochloric acid, carry out the procedure below;

Pipette 25cm3 of the sodium carbonate solution into a conical flask.

Titrate it using hydrochloric acid of unknown concentration from a burette

Add a few drops of methyl orange indicator, and titrate until the colour changes slightly

Record the volume of the acid used.

The following are sample results from the titrations

Titrations	Pilot	I	II	III
Final burette reading	21.05	20.25	40.75	20.55
Initial burette reading.	0.00	0.00	20.25	0.00
Volume of HCl	21.05	20.25	20.50	20.55

To calculate the concentration of the acid, proceed as follows;

Step 1. Calculate the average volume of the acid used.

Average titre- 20.5 +20.55 = 20.53cm3

Thus 20.53cm3 of the acid reacts with 25cm3 of sodium carbonate.

Step 2. Calculate the number of moles of sodium carbonate used.

Molarity =

Number of moles = volume of solution X molarity

Since 25 cm3 of 0.098M sodium carbonate solution was used in the reaction, the number of moles used is:

= 0.00245 moles.

Step3. Write a balanced chemical equation for the reaction.

Na2CO3(aq) + 2HC1(aq) → 2NaCl(aq) + CO2(g) + H20(1)

Ionic equation:

CO32-(aq) + 2H+(aq) →3 CO2(g) + H20(1)

Step4. Calculate the number of moles of hydrochloric acid used from the mole ratio.

Both the ionic and full formulae equations in (3) give the same mole ratio.

From the equation, the mole ratio of Na2CO3 :HC1 is 1 : 2.

Therefore, the number of moles of hydrochloric acid is

2 x 0.00245 mol = 0.0049 mol.

Step 5. Work out the molar concentration of the hydrochloric acid.

The average volume of HCl used in the experiment was 20.53 cm3.

This means that 20.53 cm3 of hydrochloric acid contains 0.0049 mol of the acid. Therefore,

1000cm3 contains

The concentration of the acid is 0.239 mol dm-3 or 0.239M.

Alternatively,

Where:

CA is the molar concentration of the acid, VA is the volume of the acid used,

nA is the number of moles of the acid used, CB is the molar concentration of the base, VB is the volume of the base used, and

nB is the number of moles of the base used.

Therefore,

C A =

Where, Na = 0.0049mol

CB = 0.098 mol/dm3

VB = 25 cm3

nB = 0.00245 mol, and

VA = 20.53 cm3

CA =

= 0.239mol/dm3

How to determine the relative atomic mass of unknown elements in an acid or a base

The atomic mass of an unknown element in an acid or alkali can be determined by titrating the acid or the alkali solution with a standard solution of an acid or a base. For example, you are given a metal carbonate X2CO3, Whose element X is an unknown.

The following titration process can be carried out to determine

Make a solution of sodium carbonate by dissolving 2.21g of the carbonate in distilled water to make 250cm3 solution.

Fill burette with 0.25M hydrochloric acid and note the burette reading.

Pipette 25cm3 of the carbonate into conical flask and add three drops of methyl indicator

Carry out titration until colour change appears, take the reading on the burette.

Read the volume of the acid used.

Read the procedure for three times and record your values on the table below.

Sample results from above experiment may be as follows;

Titrations	pilot	I	II	III
Final burette reading	17.00	16.10	32.00	46.10
Initial burette reading	0.00	0.00	16.10	30.00
Volume of HCl (cm3)	17.00	16.10	15.90	15.90

To determine the relative atomic mass of X, follow the steps below;

Step 1. Calculate the average volume of the acid used;

The average volume is determined by averaging the three values because they and within ±0.2 cm3 of each other.

= 16.03 cm3.

Therefore, the average volume required to react with 25cm3 of metal carbonate is 16.03cm3.

Step2. Calculate the number of moles of the hydrochloric acid used:

The molarity of the acid is 0.25M. This mean that 1000 cm3 of HC1 contain 0.25 moles of the acid. Therefore, 16.03 cm3 contains

Step 4. Write a balanced chemical equation for the reaction to get the mole ratio:

The equation for the reaction is:

X,CO3(aq) + 2HC1(aq)→ 2XC1(aq) + CO2,(g)

Step 5. Determine the number of moles of the metal carbonate used:

The mole ratio of the metal carbonate to the acid is 1 : 2. Therefore, the number of moles of the metal carbonate used in the reaction is;

0.00mol/2 = 0.002 moles

Step 6. Work out the molar concentration of the metal carbonate solution.

Since, 0.002 moles of the metal carbonate are contained in 25 cm3, 1000 cm3 contain

The molarity of the carbonate solution is 0.08M.

Step7. Calculate the mass of the metal carbonate in one litre of solution:

Since, 250 cm3 of the solution contains 2.12 g of the metal carbonate, 1000 cm3 of the solution contains;

Using the molarity of the solution and the mass of the- metal carbonate per litre of solution, work out the relative molecular mass of the metal carbonate.

The concentration of the carbonate solution in grams per litre is 8.48 g/dm3.

Molarity =

Therefore, molar mass =

The relative molecular mass of the carbonate is 106 g.

Calculate the relative atomic mass of the metal based on the formula of the carbonate:

The formula of the metal carbonate is X2CO3. The relative atomic mass of carbon and oxygen are 12 and 16 respectively.

Therefore, 106 = 2X + 12 + (16 x 3) 106 = 2X + 60

46 = 2X

X = 23

The relative atomic mas,/of the metal in the carbonate is 23.

Determining the percentage purity of a substance.

The percentage purity of a substance is the percentage of the pure substance in a sample.

That is, percentage purity =

For example, if a sample of sodium carbonate with a mass of 10.6 g is analyzed and found to contain 7.8 g of pure sodium carbonate, the percentage purity of the sample is:

This means that 75% of the mass of the sample is sodium hydroxide. The remaining 25% of the mass is made up of other substances.

Titration reactions are used to determine the percentage purity of common laboratory reagents.

Activity 2.8

Determination of percentage purity of an impure sample of sodium carbonate.

The percentage impurity in a solid of sodium carbonate can be determined as follows;

Dissolve about 6.63 g of sodium carbonate in water and make the solution to 50cm3

Titrate it with 0.5m HCl using methyl orange indicator and record volume used.

Carry out three successive titrations and record your readings.

Sample results.

Titration	pilot	I	II	III
Final reading	16.00	16.20	32.45	16.10
Initial reading	0.00	0.00	16.20	0.00
Volume used	16.00	16.20	16.5	16.10

Carry out the following steps to determine percentage purity.

Calculate average volume used

Calculate the number of moles of the acid

If 1000cm3 contains 0.5moles of the acid, 16.18cm3 will contain;

0.5molx16.18cm3/1000 = 0.008moles.

c. Write a balanced chemical equation

Na2CO3(aq) + 2HC1(aq)→ 2NaC1(aq) + CO2,(g)

Thus mole ratio is; 1.2

Moles of sodium carbonate used will be,

0.008/2 = 0.004 moles

f. Calculate the mass of sodium carbonate used in the reaction;

mass = number of moles x molar mass

= 0.004 x 106 = 0.424g

g. The percentage purity =25cm3 contains 0.424g of carbonate, 250cm3, will thus contain

250cm3x0.424g/25cm3= 4.24g

So, in 6.63g of sodium carbonate, only 4.24g is pure, percentage purity will be given by;

= 4.24/6.63x100 = 64%

Thus only 64% is sodium carbonate.

Determination of water of crystallization in a substance.

Water of crystallization is the water that is bound within the crystals of a substances. This water can be removed by heating.

A hydrated carbonate has water of crystallization and has a formula Na2CO3.XH2O. The following procedure can be done to establish the amount of water of crystallization;

Fill the burette with HCl up to zero mark

Pipette 25cm3 of sodium carbonate on conical flask

Titrate careful and record you readings on the table as shown below

Titration	pilot	I	II	III
Final reading	28.95	28.40	28.45	38.50
Initial reading	0.00	0.00	00.00	10.00
Volume used	28.95	28.40	28.45	28.50

The following procedure is used to determine water of crystallization

Calculate average volume used

Find moles of acid used.

= 28.45/1000 x0.24mol/dm3 = 0.0068 moles

Write a balanced chemical equation and write moles ratio

Na2CO3(aq) + 2HC1(aq)→ 2NaC1(aq) + CO2,(g). mole ratio = 1:2

Moles of the carbonate will be;

0.0068/2 = 0.0034moles

5. Determine the molar concentration of the carbonate;

Molarity of te carbonate will be;= number of moles/volume

= 0.0034mol/0.025dm3

= 0.136 mol/dm3

The molar concentration of te carbonate solution will be;

0.136mol/dm3= 0.136M

6. Calculate the RMM of the carbonate.

Molar mass = mass per litre/molarity

38.9g/dm3/0.136mol/dm3

286g/mol

7. Determine the number of moles of water of crystallization as follows;

Na2CO3.XH2O = 2Na+C+3O+X(2H+O)=286

=(2X23)+12+(3X16)X(2+1)+]=286

=106+18x=286

X=10

Therefore there 10 molecules of water of crystallization. The formula of hydrated sodium carbonate will be Na2CO3.10H20

PRACTICAL APPLICATION OF VOLUMETRIC ANALYSIS

Volumetric analysis is used in different fields. It is an important technique used by chemists. The following are areas where volumetric analysis is used;

Medicine

In medicine, titration is used to determine the concentration of a virus or a bacterium in a blood sample.

The amount of acid which can be neutralized by an antacid tablet is also determined using volumetric analysis.

Industries

The amount of acetic acid in vinegar is determined by volumetric analysis. Table vinegar should have a minimum of 5% acetic acid content. This is established by titrating a known volume of vinegar with a standard solution of a base.

The analysis of the acidity of fruit juices is done volumetrically. This is done by titrating a known volume of the fruit juice with a standard solution of a base.

The percentage of iron in water can be found by titratin with standard potassium manganate (VII) solution

The analysis of household ammonia is also carried out volumetrically. A known volume of the ammonia is titrated with a standard solution of an acid.

Analysis of water

Volumetric analysis is also used in the determining the hardness of a water sample.

Used also to determine mineral content in bottled water

Agriculture

The composition of different substances contained in the soil can be determined by carrying out titrations in the laboratory. These substances include nitrogen, potassium and phosphorous.

Other applications

Other commercial applications of volumetric analysis include:

determination of the percentage of iron content in an iron ore

determination of salt content in brine

Determination of the percentage mass of copper in a copper salt.

SUMMARY

Volumetric analysis is used to determine accurate volumes of reacting solutions

If the volumes of two solutions that react exactly with each other are measured for example an acid and an alkali, the concentration of one can be calculated if the concentration of the other is known.

The volume of solution is usually given in litres or dm3 while the concentration is given in moles per litre.

A standard solution is one in which the concentration and molar concentration are exactly known.

The procedure used to determine the concentration of a solution is called standardization

There are two types of standards, primary standard and secondary standards.

Primary standards are prepared by dissolving a known amount of a primary standard in a specific volume of a known liquid.

Secondary standard solutions are prepared by standardizing using primary standards.

A laboratory procedure used to determine the concentration of unknown solution using a standard solution is called titration.

Titration can also be used to;

Determine the relative atomic masses of unknown elements

Determine the percentage purity of unknown substance

Determine the amount of water of crystallization in a substance

Volumetric analysis is used in various fields such as medicine, industry, metal extraction and manufacturing.

END OF TOPIC QUESTIONS.

What mass of potassium chloride is needed to prepare 250cm3 of 0.235M solution?

4.38g

9.23g

15.60g

31.3g

1.0g

Which of the following laboratory equipments can accurately measure volume of a titre that has reacted completely with an aliquot?

Pipette

Burette

Measuring cylinder

Beaker

What will be the molarity of a solution which contains 26.5g of anhydrous sodium carbonate in 5dm3 of solution?

0.05 M

0.25 M

5.30 M

0.025 M

0.50 M

What volume of 0.2M H2SO4 is required to neutralize completely 25cm3 of 0.05M KOH?

0.656cm3

6.125cm3

3.225cm

3.125cm3

(v) The mass of sodium hydroxide contained in 25cm3 of 0.1M is .

0.5, (B) 2.85g (C) 25.0g (D) 0.1g (E) 25g

A solution of sodium carbonate was prepared in order to get a 2M solution. 200cm3 of this solution was used in a titration experiment. The number of moles present in 200cm3 of 2M solution used in the titration will be;

4.0 B. 0.04 C. 0.40 D. 0.045

Which indicator is suitably used for the titration of weak acid and a strong base?

Methyl orange

Phenolphthalein

Litmus

Bromothymol blue

PART II

Given that in the chemistry national four examination, chemistry practical (2A), the following data were obtained by titrating ----0.1M HCl with 5.3g/dm3 of X2 C03 using methyl orange as an indicator. Pipette used was 20cm3

Table of results

Burette readings.

Titration number	Pilot	1	2	3
Final reading (cm3)	20.00	40.00	20.00	40.2
Initial reading (cm3)	0.00	20.00	0.00	20.1
Volume used (cm3)

Questions.

Complete the table above

Calculate titre mean volume

Write a balance chemical reaction between X2CO3 and HCL

Calculate molarity of X 2CO3

Identify and name element X

What do we mean by volumetric analysis?

State the use of each of the following apparatus used in the volumetric analysis:

Volumetric flask

Burette

Pipette

Measuring cylinder

Weighing balance

Dropper.

An experiment was carried out to find the value of X in Na2CO3.XH2 . O.12.4g of Na2CO3.XH2O was dissolved in water to make 0.5dm3 of solution. 25cm3 of this solution was transferred to a conical flask and few drops of methlyorange(MO) was added.The solution in a conical flask was then titrated against 0.25M HCl and 20cm3 of acid were required to reach the end point.

i) write balanced chemical equation for the reaction

ii) Calculate concentration of Na2CO3.XH2O in mol/dm3.

iii) Calculate the value of X.

6cm3 HCL required 22 cm3 Of 0.2 M NaOH for complete neutralization. Calculate the concentration of the acid in moles per time.

Solution Q was made by dissolving 2.65g of a X2CO3 in water and diluting it to 250cm3. 25 cm3 of P was titrated with 0.25 MHCL solution. The results obtained were tabulated as shown.

	I	II	IV
Final burette reading (cm3)	18.90	37.9	19.0
Initial burrete reading (cm3)	0.00	18.90	0.00
Volume of HCL (Cm3)	18.90	19.00	19.10

Calculate the average volume of HCL acid used.

Calculate the concentration of solution Q in:

I) Mol /dm3

II) g/ dm3

Calculate the molarities of the following solutions:

10.og of sodium hydroxide, NaOH, in 250cm3 of solution

3.5g of iron (II) Sulphate FeSO4 in 100 Cm3

Calculate the molarities of the following solutions:

10.0 g sodium hydroxide, NaOH, in 250 Cm3 of solution

3.5g of iron (II) sulphate FeSO4, in 100cm3 solution.

Calculate the number of moles in:

20cm3 of 0.1M Nitric (v) acid.

150cm3 of 0.5 Sodium carbonate.

(a) Explain the meaning of the following.

Titration

Molar solution

(b)Given that 5.6g of KOH were dissolved in 1dm3 .25cm3 of this solution ( KOH) were titrated with 25cm3 of hydrochloric acid. Calculate;

(i) Molarity of acid

(ii) Concentration of KOH

Name suitable indicator for this titration.

(a) Define

End point of neutralization

Indicator

Standard solution

Secondary standard solution

(b) A solution of sulphuric acid is made by dissolving 5.2g of its mass containing some

impurities to make 1 dm3 of solution.

If 25cm3 of this solution required exactly 25cm3 of 0.1M NaOH for complete

neutralization during the titration process. Calculate the percentage purity of sulphuric acid.

Define the following terms:

Standard solution

End points

Titration

190 cm3 of the 0.4M lithium carbonate solution was neutralized by 16.0Cm3 of 1MH2SO4. Write an equation for the reaction.

In an experiment, 22.00cm3 of 1.0 NaOH Was neutralized by 170.0 cm3 of nitric (v) Acid. What is the concentration of nitric (v) acid in mol/ dm3.

Calculate the molarities of the following solutions:

25.0g of sodium hydroxide, NaOH in 200cm3 of solution.

3.5 g of iron (II) sulphate, FeSO4, in 250cm3 of solution.

Calculate the volume of

0.1M Na,CO3 solution that neutralizes compeletely 2400cm3 of 0.2M H2SO4 Solution.

2M nitric (V) acid required to neutralize 25.00 cm3 of 0.5M NaOH.

Xg of Sodium hydroxide were dissolved in distilled water to make 100cm3 of solution . 50 cm3 of the solution required 50cm3.

2M nitric (V) acid for complete neutralization. Caliculate neutralization. Calculate the mass (x) of Sodium hydroxide dissolved ( Na ═ 23 O ═ 16 H ═ 1)

Calculate the amount of calcium carbonate that would remain in 17.0g of calcium carbonate were reacted with 0.25 Moles of hydrochloric acid. The equation for the equation for the reaction is CaCO3(g) + 2HCL(aq) → CaCl2(aq) + CO2 + H2O(I)

(C ═ 12 O ═ 16 Ca ═ 40).

0.88gcm-3. Calculate the morality of the ammonia solution.

How would you prepare 0.15 L of 0.5 M NaOH, Starting with 6.0M Solution.

How would you prepare 0.75 litre of a 0.5M NaoH Solution, starting with solid sodium hydroxide and water.

Calculate the volume of 1M sulphuric (vi) acid need to completely neutralize 100cm3 of sodium hydroxide 0.2M in concentration.

The reaction below shows, reaction between calcium hydroxide and nitric (iv) acid.

Ca(OH)2(S) + 2HNO3(aq) → Ca(NO3)2(aq) + 2H2O(l)

If 20 Cm3 of 0.2 M nitric (v) acid was used, calculate the volume of Ca(OH)2 used, given its morality is 0.1M

3M nitric acid reacted with 0.5g copper powder as shown below

CU(S) + 2HNO3 (aq) → Cu (NO3) 2 + H2O

Given (CU = 63.5)

Calculate the volume of 3M nitric acid that reacted with copper.

Calculate the mass of nitrogen present in 25kg bag of ammonium phosphates (NH4)2 HPO4.

(N = 14, H= 1, P = 31, O = 16)

When 1.08g of aluminium foil were in a stream of chlorine gas, the mass of product formed was 3.47g. calculate the

Maximum mass of product formed of chlorine was excess (AL 27, CL = 35.5)

Percentage yield of the product formed.

15.0 cm3 of ethanoic acid (CH3COOH) was dissolved in water to make 500cm3 of solution. Calculate the concentration of the solution in moles per litre.

(C= 12.0 H = 10, O = 16.0) (Density = 1.05g/cm3 of ethanoic acid)

The relative formula mass of a mass of hydrocarbon is 58. Draw and name two possible structure of the hydrocarbon

An alcohol has this composition

H= 13.5%

O2 = 64.9 %

Determine the structure of alcohol if empirical formula and molecular are the same.

Draw the structure of alcohol if empirical formula and molecular formula are the same.

6.84g of aluminum sulphate were dissolved in 150cm3 of water calcululate molar concentration of the sulphate ions in the solution. In of aluminium sulphate is 342).

Work out the simplest formula of a compound. Which has the follows composition

Magnesium 9.8g

Sulphur 13.0g

Oxygen 26.0g

Water of crystallization 51.2g

150g of Calcium carbonate were reacted with 0.2 moles of hydrochloric acid.

Write equation of reaction

Calculate amount of calcium carbonate that will remain unreached.

0.7 g of gas × occupies 560 cm3 at s.t.p. determine its relative molecular mass.

Molar gas vol = 22.4L

A certain carbonate aCO3 reacts with dilute hydrochloric acid according to the equation given below.

ACO3(s) + 2HCl (aq) → al2aq + CO (g) + H2O (l)

If 1g of the carbonate reacts completely with 20cm3 of the IM HCL.Calculate the relative atomic mass of G. (C = 12.0, O = 16.0)

When a hydrocarbon was completely bunted in oxygen, 4.2g of carbon dioxide has1.71g of water were formed. Determine the empirical formula of late hydrocarbon. ( H= 1.0, C = 12.0 , O = 16.0 )

(a) What is titration?

(b) 25cm3 of impure sulphuric acid contains 5.2g/dm3 reacted with 23cm3 of sodium hydroxide solution made by dissolving 4.0g of NaOH in distilled water to make 1.0dm3 solution

(i) Purity of the acid and (ii) its impurity

(d) Mention two areas of application of volumetric analysis

20.0cm3 of a 1.4M, Dibasic acid H2X solution was titrated against 1.0M NaOH solutionuntil neutralization was complete.

Write down the equation for the reaction.

Calculate the volume of sodium hydroxide solution used.

SAMPLE NECTA 2009.

QN. 12.(a) Define the following terms:

Standard solution

Equivalent point of a titration

Basicity of an acid

(b) 25cm3 of a solution containing 0.196g of a metal hydroxide YOH, were neutralized by 35cm3 of a 0.1M hydrochloric acid solution.

Write down a balanced chemical equation for the reaction.

Calculate the molarity of the hydroxide solution.

Calculate the relative molecular mass of YOH.

Hydrochloric acid versus ammonia solution

Acetic acid versus sodium hydroxide

2010.

QNS. 12. Read the following information carefully then answer the questions that follow:

(a) (i) What piece of apparatus should be used to measure out accurately 25cm3 of sodium hydroxide solution?

(ii) What colour was the solution in the flask at the start of the titration?

(iii) What colour did it turn when the alkali had been neutralized?

(b) (i) Was the acid more concentrated or less concentrated than the alkali? Give reasons for your answer.

Name the salt formed in the neutralization

Write an equation for the reaction.

2013.

QNS(a) 25cm3 of 0.1M HCI were neutralized by 23cm3 of sodium hydroxide solution. Calculate the concentration of the alkali in grams per litre.

2013.

QNS. 4. (a) 20cm3 of a solution containing 7g dm-3 of sodium hydroxide were exactly neutralized by 25cm3 of 0.10 M hydrochloric acid. Calculate the concentration of sodium hydroxide in moles per dm3.

(b) Give two examples in each of the following solutions:

Gaseous solution

Solid solution

TOPIC : 7 IONIC THEORY AND ELECTROLYSIS

P CLASS="western" ALIGN=CENTER STYLE="margin-bottom: 0.11in">CHAPTER 7

IONIC THEORY AND ELECTROLYSIS.

KEY CONCEPTS AND TERMS.

Conductors- are solids that can conduct electric current.

Insulators- these are solids that cannot be able to conduct electric current. They are also called insulators.

Poor conductors- are substances that only a small amount of electric current to pass through.

All metals are good conductors of electricity. Non metals do not conduct electricity except carbon( graphite). The metals conduct electricity by used of electrons which are able to move freely.

Electrolyte- is a substance which dissociate into free ions when in solution or in molten state, thus allowing electric current to pass. Electrolytes includes, ionic compounds such as sodium chloride, salts, bases and acids. Electrolytes conducts electricity with help of free ions.

Non- electrolytes- is a solution or molten compound which does not conduct electricity.

Electrolysis- is the process by which an electric current is passed through an electrolyte to cause a chemical reaction.

Electrolytic cell- is an apparatus in which electrolysis is carried out

Strong electrolyte- this is a type of electrolyte in which ions are fully ionized

Weak electrolyte- this is an electrolytic solution in which ions are partially ionized

Electrochemical series- this is the arrangement of elements depending on their ease in gaining electrons.

An anode- this is a terminal where electrons leave the electrolyte.

Cathode- this is the terminal where electrons enter the electrolyte

Faradays first law of electrolysis- this states that;’’the mass of a substance produced at the electrodes during electrolysis is proportional to the quantity of electricity passed. In actual sence, the number of moles of a substance produced at an electrode during electrolysis depends on the quantity of electricity passing through the electrolyte, amount of time taken, and the charge on the ion.

Coulomb- this is the quantity of electric charge that passes through a given point in a circuit when a current of 1 ampere flows for 1 second.

Electrochemical equivalent (Z)- Is the mass of the substance discharged when 1 coulomb of electricity is passed through electrolyte.

One faraday- this is the quantity of electric charge carried by one mole of electrons.

Faradays second law states that; ‘’when the same quantity of electricity is passed through solutions of different electrolytes, the mass of the substances liberated or deposited at the electrodes is directly proportional to the chemical equivalent of the substance.

Electroplating- this is the process by which a metal is coated with another one by passing an electric current through electrolyte.

Ionic theory.

This theory, states that ionic compounds are up of positive and negative ions. The positive ions are called cations, while the negative ions are called anions.

When ionic compounds dissolve in water, they dissociate to form free ions.

NaCl(s)

Na+ + Cl-

Note that most ionic compounds are solid at room temperature.

When substances made of ionic compounds are dissolved in water, they dissolve forming ionic solutions, which contains ions that conduct electricity. Since the ionic compound is dissolved in water, it will also contain ions of water. For example, in a solution of calcium chloride, there will be Ca ions, chloride ions hydrogen ions and hydroxyl ions.

Electrolysis.

This is the process by which an electric current passes through electrolyte causing a chemical reaction. The solution can be molten or in solution. The role of water in any electrolyte is to make the ions mobile and thus enable them carry electric current.

Electrolytic cells.

The process of electrolysis takes place in special cells called electrolytic cells. An electrolytic cell consists of the following;

Battery- It is shown by the symbol where the longer and thinner stroke represents the positive terminal while the short and thick stroke represents the negative terminal.

Electrodes- there are two types, the positive electrodes where electrons leave the electrolyte is called anode. It is attached to the positive terminal of the battery. The other electrode is called the cathode and is attached to the neagative terminal of the battery. This is the terminal where electrons enter the electrolyte.

Electrolytes- this is the solution in which the electrodes are dipped. It should be molten or liquid in nature.

During electrolysis, oxidation takes place at the anode while reduction takes place at the cathode.

The electrodes used should be made of inert material so that they do not react with the solution.

Weak and strong electrolytes.

Weak electrolytes- These are electrolytes which donot dissociate completely in solution or in liquid state, but do dissociate partially. They are not very efficient in electric conduction because the ions are few in number. Examples includes, Ammonium solution, Ethanoic acid and Methanoic acids.

Strong electrolytes- these are electrolytes which dissociate completely when dissolved in water. Most acids and salts are strong electrolytes.

Weak electrolytes-

NH4Cl

NH4+ + Cl-

CH3COOH

CH3COO- + H+

HCOOH

COOH- + H+

Strong electrolytes.

Sulphuric acids.

H2SO4(aq)

H+ + SO42-

Hydrochloric acid.

HCl(aq)

H+ + Cl-

When ionic compounds are dissolved in organic compounds, they remain in molecular form and they do not conduct electricity. Organic solvents can not dissociate into free ions. This is why hydrochloric acid does not conduct electricity in benzene.

Concentration and strength of an acid.

Strength of an electrolyte- this is the measure of the extend of dissociation of an acid. The more the electrolyte dissociates, the stronger it is.

Concentration- this is the measure of the number of moles of a solute in a given volume of solution..

Electrolytes with high concentration can be weak and vice versa. They can also have the same strength but different concentration.

Mechanism of electrolysis.

This is the movement of electrons to the electrodes. The positive ions or Cations moves to the cathode( negative terminal), while the negative ions(anions) moves to the anode, or the positive terminal.

Anions such as sulphate, nitrate, chlorides, bromides, iodide and hydroxyl moves to the anode. Cations such as potassium, sodium, calcium, magnesium, alluminium zinc iron lead hydrogen copper and silver moves to the cathode.

Preferential discharge of ions during electrolysis

When more than one type of electrodes migrate to the electrodes, one of them is easily discharged compared to others.

This selectiveness in discharge is called preferential discharge. In addition to the ions of the solute, an aqueous solution of an electrolyte contains hydrogen ions (H+) and hydroxide ions (OH-) from water (the solvent).

The hydrogen ions together with the cations of the electrolyte migrate to the cathode while the hydroxide ions and the anions of the electrolyte migrate to the anode.

At the electrodes, some of the ions gain or lose electrons. This has the effect of discharging the ions. Only one type of ions is discharged at each electrode.

The discharge depends on:

The position of ion in electrochemical series.

The concentration of the competing ions in the solutions.

The nature of the electrode.

1. The position of the ion in electrochemical series.

For cation and hydrogen, the ease of discharge of an ion is determined by the position of the element in the electrochemical series.

The higher the ions in the electrochemical series, the more difficult it is to be discharged. The less reactive the element, the easier it is to be discharged.

The lower the ion in the series, the easier it is to be discharged. For example, during the electrolysis of copper sulphate solution, ions at the anode will be sulphate ions (SO42-) and hydroxyl ion (OH-).

The hydroxyl ion will be preferentially discharged at the anode. At the cathode we shall have Copper ion (Cu2+) and Hydrogen ions (H+), the hydrogen ions shall be preferentially discharged at the cathode.

This is due to their position in the electrochemical series.

The concentration of the ions

A higher concentration of ions tends to favour its discharge. This factor is very important for anions. In most cases it affects the products formed at the anode.

The higher the concentration of the ions, the more easy for the ion to discharge. A good example is in the electrolysis of concentrated sodium chloride ( brine), where the chloride ions are preferentially discharged.

But dilute sodium chloride will yield the hydroxyl ions at the anode.

Nature of electrodes

The nature of electrode used influence the choice of ion for discharge. Consider the electrolysis of NaC1 solution, using platinum or mercury cathode, different products will be formed.

When platinum cathode is used, H+ ions are discharged in preference to Na+ to give hydrogen gas.

2H+(aq) + 2e- ― H2(g)

When mercury cathode is used, Na+ is discharged in preference to H+ ions. This is because the discharge of Na+ requires less energy than the discharge of H. The product at the cathode is sodium amalgam.

Electrolysis of selected substances.

Products formed at the electrodes depends on the ions discharged at the electrodes.

The ions discharged depends on three factors, the position of the ion on the electrochemical series, the concentration of the ion and the nature of electrodes.

The following section deals with the electrolysis of selected substances

Electrolysis of sodium chloride.

Sodium chloride contains sodium ions(Na+), chloride ions Cl- all from sodium chloride and OH- and H+ from water.

Figure 7.1 electrolysis of sodium chloride

Reaction at the anode.

A colourless gas which relights a glowing splint is collected.

OH is discharged in preference to C1-(aq) because OH-(aq) is lower in electrochemical series than C1-(aq).

40H-(aq) → 02(g) + 2H20(1) + 4e-

Product; 1 volume of oxygen

At the cathode

A colourless gas which burns with a "pop sound" is collected

H+(aq) is reduced in preference to Na+(aq) because H+(aq) is lower than Na+(aq) in the electrochemical series.

2H+(aq) + 2e- → H2(g)

Product; 1 volume of hydrogen

Overall ionic equation at the cathode

4H+(aq) + 4e- → 2H2(g)

The electrolysis of dilute sulphuric acid.

The solution contains hydrogen ions and sulphate ions from sulphuric acid, and hydrogen ions and hydroxide ions from water.

Figure 7.2 dilute sulphuric acid

Reaction at the anode

A colourless gas which relights a glowing splint is collected. Both OH-(aq) and SQ42-(aq) move to the anode, but the OH-(aq) gets discharged in preference to SO42-(aq).

40H-(aq) → 2H20(1) + 02(g) + 4e-

At the cathode

A colourless gas which burns with a "pop sound" is collected.

The H+(aq) are the only cations in the solution. They move to the cathode where they are discharged.

2H+ (aq) + 2e → H2(g)

Product; 2 volumes of oxygen

The 4 electrons from the anode are taken up by H+ ions giving 2 moles of hydrogen gas.

4H+ (aq) + 4e → 2H2(g)

Electrolysis of sodium hydroxide solution

The solution contains Na+(aq), OH-(aq), H+(aq) ions.

At the anode

A colourless gas that relights a glowing splint is collected. OH-(aq) ions move to the anode and get discharged.

40H-(aq) → 2H20(1) + 02(g) + 4e-

At the cathode

A colourless gas which burns with a "pop sound" is collected.

4H+(aq) + 4e→ 2H2(g)

Electrolysis of copper (II) sulphate solution

Determine the products for the electrolysis of copper (II) sulphate solution when:

Carbon electrodes are used

Copper electrodes are used

The ions present in copper (II) sulphate solution are:

Cu2+(aq), SO42-(aq), H+(aq) and OH-(aq)

Using carbon electrodes

Carbon electrodes are inert. Therefore they do not dissociate into the electrolyte.

The deep blue colour of the solution slowly fades away as more Cu2+ ions are deposited at the cathode as copper solid. The solution finally turns colourless.

Figure 7.3 electrolysis of copper sulphate using copper electrodes.

At the anode

Ions present are: SO42-(aq) and OH-(aq) OH- ions will be preferentially discharged 40H-(aq) 2H20(1) + 02(g) + 4e-

At the cathode

Ions present are: Cu2+(aq) and H+(aq) Cu2+ ions will be preferentially discharged

2Cu(aq) + 4e-→ 2Cu(s)

Using copper electrodes

When copper electrodes are used in the electrolysis of copper(II) sulphate the deep blue colour of the solution does not fade away even though the copper solid are being formed at the cathode. This is because the anode keeps dissolving into solution. With time, the size of the anode reduces as it is eaten away into the solution.

At the cathode

The ions present are: Cu2+(aq) and H+(aq) ions

Cu2+(aq) ions are preferentially discharged.

Cu2+(aq) + 2e- → Cu(s)

Products; copper

At the anode

The ions present are: SO42-(aq) and 0H-(aq)

None of the ions is discharged. Instead the copper electrode goes into solution

Cu(s) → Cu2+(aq) + 2e-

Electrolysis of molten sodium chloride

The molten sodium chloride has Na+ ions and Cl- ions.

NaCI(1) →Na+ + Cl-

During electrolysis Na+ ions move towards the cathode while the Cl- ions move towards the anode.

Each sodium ion takes up one electron from the cathode and becomes sodium metal. At the anode, each Cl- ion gives up an electron to the anode and becomes a chlorine atom.

Since chlorine atom is not stable, two chlorine atoms combine to form a chlorine molecule (C12).

The reactions may be expressed as,

At the anode

Cl- ion is oxidized to Cl atom which combines with another Cl atom to form a molecule.

Cl-→ Cl + e-

2C1- → C12(g) + 2e-

At the cathode Na+ + e- → Na(s)

Since two electrons are released to produce one mole of chlorine, an equal number of electrons shall be gained by the sodium ions. Thus the equation at the anode shall be;

2Na+ + 2e →2Na(s)

This implies that each time one mole of chlorine is produced at the anode two

moles of sodium metal are produced at the cathode.

Electrolysis of water

Pure water is a poor conductor of electricity. However, when a few drops of sulphuric(VI) acid are added to water, becomes an electrolyte.

H20(1) → H+(aq) + OH-(aq)

At the anode

OH-(aq) → OH(g) + e-

OH- is not a stable compound, thus it form a stable compound more hydroxyl ions need to undergo oxidation.

OH-(aq)→ OH + e-

40H-(aq)→ 2H20(1) + 02(g) + 4e-

At the cathode

H+ ions moves to the cathode and gain au electron to from H atom.

H+(aq) + e → H(g)

Since hydrogen is a diatomic gas, the atom combine to form a hydrogen molecule.

H(g) + H(g) → H2(g)

The 4e- from the anode are taken up by H1 ions giving 2 moles of hydrogen gas

4H+(aq) + 4e- → 2H2(g)

Laws of electrolysis.

Michael faraday observed that there was direct relationship between the amount of products formed at the electrodes and the quantity of electricity that is passed through the electrolyte. Michael summarized these ideas in what is commonly called faradays laws of electrolysis.

Faradays first law of electrolysis

The amount of any substance deposited or liberated at any electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

The amount of any substance obtained, gives the amount of chemical reaction which occurs at any electrode during electrolysis.

Thus if m gram of the substance is deposited on passing Q coulombs of electricity, then;

m α Q and m = ZQ

Where Z is the proportionality constant and is called the electrochemical equivalent.

If a current I is passed for seconds t then Q=Ixt

So, m=ZxIxt= ZIt

If I = 1 ampere and t = 1 second, then m= z

Z = Ar/V.F

Where Ar = mass of one atom of an element in g.

V = Valency

F = 96,500 Coulombs

Hence, the electrochemical equivalent of a substance may be defined as the mass of the substance deposited when 1 ampere of current is passed for 1 second, (quantity of electricity passed is equal to 1 coulomb). It can also be defined as the mass of the substance discharged when 1 coulomb of charge is passed through an electrolyte.

If m =Zxlxt

and Z = Ar

V.F

then

m can also be written as

Ar x It

V.F

Faradays constant is the quantity of electricity required to liberate 1 mole of a monovalent substance during electrolysis.

How to calculate Z of elements.

Calculate the electrochemical equivalent of hydrogen

(Ar of H = 1.008, 1F = 96500C)

Given Ar = 1.008g, Z = Ar/V.F

1F = 96500C

Z=?

V=1

= 1.008/96500

0.0000104g/C

T o verify faradays first law of electrolysis.

This can be done as follows;

An electrolytic cell made of copper sulphate solution is set up. Copper electrodes are used as both cathode and anode. The weight of both electrode is measured at the beginning of the experiment. Then the circuit is closed to allow electric current to pass. The mass of copper deposited at the anode is measured at different intervals.

Observations.

It is found that the mass of copper deposited is directly proportional to the amount of electricity passed. The more the amount of current is passed, the greater the mass deposited. This is in agreement with faradays first law of electrolysis which states that; The mass of a substance produced or dissolved at the electrode during electrolysis is proportional to the moles of electrons transferred at the electrodes

The electrochemical equivalent is the mass of a substance discharged when I coulomb of electricity is passed through electrolyte .

One coulomb is the quantity of electricity charge that passes through a given point in a circuit when a current of 1 ampere flows for 1 second.

Faradays constant

The amount of electric charge carried by one mole of electrons is known as Faradays constant.

The charge on 1 electron is 1.602 x 10-9C. The charge on 1 mole of electrons (6.022 x 1023 electrons) is therefore 1.602 x 10-9C x 6.022 x 1023

9.647 x 104 C mol- = 96 470 C

The charge on one mole of electrons is usually approximated to 96 500 C. This is the Faradays constant.

1 mole of copper is equivalent to 63.5 g of copper and the electrochemical equivalent of copper is 3.333 x 10-4 g C-. Therefore, the quantity of electricity required to deposit 1 mole (63.5 g) of copper is

63.5g/mol/0.0003333C = 190,519Cmol

The equation for the deposition of copper is: Cu2+ (aq) + 2e → Cu(s)

This means two moles of electrons are needed to deposit one mole of copper (63.5 g) atoms or 190 519 C deposit one mole of copper atoms. The quantity of electricity carried by the two moles of electrons is 190 519 C.

Therefore, one mole of electrons carries:

190519C/2 = 95,259.5

This is the same as faradays constant.

Faraday’s constant as a unit of measurement

Faraday’s constant (96 500 C) is also used as a unit of measurement called faraday and abbreviated as F.

One faraday (F) is the quantity of electric charge carried by one mole of electrons.

Generally, the amount, in moles, of atoms of an element liberated or deposited at an electrode is equal to:

amount in moles of electrons passed

number of charges on each ion discharged.

One faraday discharges one mole of silver ions or one mole of sodium ions.

Ag+ + e- → A g (s )

K + + e- → Na(s)

Two faradays are required to discharge one mole of magnesium or calcium ions:

Mg 2+ + 2e- → Mg(s)

Ca 2+ + 2e- → Ca(s)

Three faradays are needed to discharge one mole of aluminium ions:

Al 3+ + 3e → Al(s)

Therefore the number of faradays of electricity is equivalent to the electrons required to deposit one mole.

NB. This is a very important concept when carrying out calculations involving mole concepts.

Faraday second law of electrolysis.

This law states that;’’ When equal quantity of electricity is passed through solution of different electrolytes, the mass of the substance deposited or liberated at the electrodes is directly proportional to the chemical equivalent of the substances.’’

Chemical equivalents, E

Chemical equivalent of an element is the combining weight of the element.

The combining weight is the mass in grams of an element that will combine with or displaces 1g of hydrogen.

Chemical equivalent = R.AM/ Number of charge on an ion

According to faraday second law;

M1/m2 = E1/E2

Where m1 and m2 are the respective masses of substances liberated at or deposited at the electrodes while E1 and E2 are chemical equivalents of the two substances respectively.

Thus when the same quantity of electricity is passed through a number of electrolytic solutions connected in series, then the masses of the different materials are liberated at the respective electrodes are in the ratio of their chemical equivalent masses.

Chemical equivalents are usually obtained by dividing the formula mass of an element by its valence.

M = kE

Where

M; The mass liberated during electrolysis K; The constant

E; The chemical equivalent

Ar; Relative atomic mass

V; Valence of element

For first electrolyte: M1= kE1 (i)

For second electrolyte: M2 = kE2 (ii)

Then dividing equation (i) by equation (ii) we obtain

M2 – E2

To verify faraday second law.

This law can be verified by setting up apparatus as follows;

Set up two electrolytic cells, one made up of copper sulphate electrolyte while the other one made up of silver nitrate solution all of the same concentration.

The two cells are connected by a circuit with a rheostart and a battery of 6V. The electric current is allowed to flow for 35 minutes.

The mass deposited at each of the electrodes of the cell is calculated.

Sample results, it was found that the mass of copper deposited was 0.449g while that of silver deposited was 1.542g

M1/m2 = 0.449/1.542 = 0.29

The chemical equivalent of copper is 63.5/2 = 31.75

The chemical equivalent of silver = 108/1 = 108.

But E1/E2 = 31.75/108 =0.29

THUS m1/m2 =E1/E2

This is in agreement with faraday second law of electrolysis.

Applications of electrolysis.

The following are some of the industrial applications of the process of electrolysis.

Industrial purification of copper.

This is refining of copper after it has been extracted. This process uses electrolysis. In this case, impure copper is used as anode in electrolytic cell while pure copper is used as cathode. When electric current is allowed to pass through the solution of copper sulphates, the impure copper dissolves decreasing in mass while the pure copper increases in mass. The copper ions from impure copper get discharged at the cathode increasing its mass. The impurities from the anode remain deposited on the floor of the electrolytic cell.

Reaction at the cathode.

Cu2+ (aq) + 2e- → Cu(s)

Reactions at the anode

Because the copper at the anode require less energy to be oxidized, it is oxidized into copper ions

Cu(s) → Cu2+ (aq) → 2e-

Electroplating of metals.

Is the process whereby a thin layer of material is applied on another object.

It is done to prevent rusting and sometimes to improve the appearance of the material. Electroplating is the electrical deposition of one metal on another.

In electroplating, a cheaper metal is usually coated with a more expensive one.

The purpose of electroplating is to protect the metal being electroplated from corrosion or rusting and to decorate it or improve its appearance.

Electroplating is used to make ornaments such as rings, earrings, bangles and wrist watches to appear as if they are made of gold or silver.

Steel parts of cars, bicycles and some tools appear shiny because they are electroplated with chromium.

During the electroplating process, the anode is usually made of the metal to be used for plating while the object to be plated is made the cathode.

The electrolyte used should be a salt solution of the plating metal. During the electrolysis, the plating metal dissolves and is transferred to the cathode (the object being plated).

SUMMARY

Electrolytes are ionic substances that conduct electricity and get decomposed by it.

Non metal electrolytes are covalent substances that do not conduct electricity and hence are not decomposed by it

Weak electrolytes allows small amount of an electric current to pass through them hence produces few ions.

Strong electrolytes allows a large amount of an electrolyzed

Electrolytes commonly exists as solutions, or melts of salts, acids and bases

Strong electrolytes completely dissociate into ions when in solution or in molten state . most acids and bases are strong electrolytes.

Electrolysis is carried out in an electrolytic cell, whose components includes battery, electrodes, and an electrolyte.

During electrolysis, cations(positive charged particles migrate to the cathode while anions(negative ions migrate to the anode.

Discharge of ions is affected by factors such the position of ion in electrochemical series, concentration of ions and the nature of the electrodes.

Faradays laws of electrolysis relates the quantity of electricity passed to the amount of substance liberated or deposited at the electrodes.

Faradays constant is equal to 96,500C. This is the amount of electric charge carried by one mole of electrons.

Electrolysis is applied in the production ofgases, purificationof metals and in the electroplating of objects.

IONIC THEORY AND ELECTROLYSIS END TOPIC QUESTIONS.

During electrolysis the molten Aluminium oxide, 3 Faradays were needed to deposit one mole of aluminium. The number electrons of aluminium will be :

6.02 x 1023

1.806 x 1023

18.06 x 10.23

180.6 x 1023

1806 x 1023

If a steady current of 2 amperes were passed through an aqueous solution of iron deposited are the cathodes will be:-

54g

56g

0.54g

28g

0.52g

The use of electricity to decompose molten sodium chloride into its components elements is an example of:-

Electroplating

Galvanization

Hydrolysis

Precipitation

Electrolysis

The charge of one mole of elections is represented by the term;

One ampere

One Faraday

One Coulorm

One volt

One gram

An electric current of 0.2A was passed through and electrolyte for 15minutes and 35 seconds. What is the quantity of electricity will be produced.

4675C

10C

187C

200C

18.7C

During electrolysis, the mass of element liberated is proportional to:-

Quantity of electricity

Chemical equation

Isotope

Faraday constant

Valency

What is the meaning of electrolysis?

Decomposition of a chemical compound by electricity

Breaking up by using electric power

Oxidation reduction process

Passing electricity through a conductor

Which is the correct property of an electrolyte

It contains free charge carriers

It does not have free ions

It can not be decomposed by an electric current

It contains mainly undissociated molecules

Which of the following is not an electrolyte

Aqueous sulphuric acid

Solid common salt

Molten calcium chloride

Alcohol in water

When dilute sodium chloride solution is electrolysed using carbon electrodes which substance is collected at the cathode

Hydrogen

Sodium

Chloride

Oxygen

What are cations?

Cathode going ions

Anode moving ions

Negatively charged ions

Unionized particles

Which particles will be anions in the electrolysis of aqueous copper sulphate using platinum electrodes?

SO42- and oH

Cu 2+ and H+

Cu 2+ and SO4 2-

Platinum

What is the name given to quantity of substance deposited by one coulomb of electricity on the electrode?

Electrochemical equivalent

Chemical equivalent

1 mole of the substance

1 G.M.V at s.t.p.

Which one of the following is the application of electrolysis

Electroplating

Isolation of elements

Purification of metals

All of the above

What is the name of the vessel in which electrolysis process is conducted

Voltameter

Voltmeter

Avometer

Beaker

How many voltameters may be required to study faradays 2nd law of electrolysis?

One

Two

Two or more

Non of the above

Which statement leads to the principle of electroplating

The article to be plated is cleaned and made the cathode

Electrolyte to be plated is cleaned and made the cathode

The anode should produce metal irons same as those in the electrolyte

All of the above

If 400cm3 of hydrogen is collected at stp from a hydrogen voltammeter, how many grams of copper should be collected from a series connected copper voltameter?

1.134g

63.50g

31.70g

0.5724g

Which formula represents faradays 1st law of electrolysis?

A. m α Ar/Z

B. Q α Z

C. n α 1/Z

D. m α Q

(v) Where is the application of volumetric analysis not found?

A. Hospitals

B. Industries

C. TBS and TFDA

D. Home

PART II

(a) A current of 0.5A was made to follow through silver voltmeters for 30 minutes.

Calculate mass of silver deposited and equivalent weight of silver.

(b) Calculate the volume of gases that would be produced at electrodes when 0.05 mole of

Electrons are passed through a dilute solution of sulphuric acid.

Define the following terms

Electrolyte

Non electrolyte

Differentiate Between a weak electrolyte and strong electrolyte.

Study the diagram below and answer the questions that follows:

Using letter X and Y Label the anode and the cathode.

Write the ionic equations for the reactions taking place at the electrodes.

Briefly explain what will happen to the happen to the mass of copper anode.

(a) Define the following terms:-

Electrolyte

Electrolysis

(b)(i) What are the three factors which effects the discharges of ions during electrolysis?

(c ) if 9.72 x 105 C of electricity were passed through a molten chloride of metal Q suggest and identify metal Q if 120.9g of it were deposited at the cathode, metal Q is divalent element.

A current of 8A was passed through a solution of an electrolyte for 600 seconds.

Calculate the number of coulombs of electricity used.

How many Faraday were used?

State the applications of electrolysis in industries.

(a) By the help of a well labeled circuit diagram explain the electrolysis of a dilute

Solution of sodium chloride using graphitic carbon electrodes.

(b) Is the above electrolysis having special name? If yes explain.

(c) Is there any further reaction at the electrodes if the electrolysis is continued near to dryness of the electrolyte? Explain your answer.

(a) Explain the meaning of “an ionic equation”

(b) Write an ionic equation for each of the following:

Carbon dioxide gas dissolves in an aqueous solution of potassium hydroxide

The reaction between magnesium and dilute hydrochloric acid.

In an electrolysis of a salt solution of metal W, a current of 0.8A was passed for 40 minutes and the amount of metal deposited was 0.56g (W= 56)

Calculate the charge carried by the metal W ion reaction

Classify the substance listed below in the table provided.

Sugar solution, solid calcium carbonate, granites, wax, sulphur, aqueous sodium chloride, distilled, water, paraffin, concentrated sulphuric (VI) acid, sodium chloride crystals, copper wire , molten copper (II) bromide.

Conductor	Non - conductor

What is the meaning of the following terms,

Cation

Anion

Cathode

Anode

A. Define the term electrolysis.

B. Name the particles responsible for the electrical conduction in

i. Copper metal

ii. Molten Lead (II) iodine

iii. Sodium Chloride Solution

The electrolysis of dilute sulphuric acidic also called electrolysis of water. Explain.

Describe the factors that determine the products found at the electrodes during electrolysis.

Magnesium sulphate is electrolyzed using a current of30.A for 30 seconds. Calculate the charge of magnesium ion if 1.12g of magnesium is deposited.

23.4g of copper is deposited within 40 mins. Calculate the current passed.

Sodium hydroxide is a stronger base than ammonium solution, though the solution has the same concentration. Clearly explain the differences between strength and concentration of an electrolyte.

4.165g of an element z was deposited from its salt solution by passing of 13510C of electricity.

Calculate the quantity of electricity that dissolve one mole of Z ( Z= 119)

Determine the charge of ion of element Z

Explain the following observations

Lead (II) iodide does not conduct electricity in solid state but conducts in molten state.

Sugar solution does not conduct electricity while salt solution does.

A copper spoon was electroplated using silver.

Calculate the mass of silver deposited on the spoon if a current of 0.05A Was passed for 3 hours. ( 1F ═ 96.500C, Ag ═ 108)

Calculate the quantity of electrolysis used.

What is electrolysis of dilute copper sulphate solution using copper cathode and copper anode.

Match the items in list A with those from list B by writing the letter of the correct response from List B besides the item number in List A.

LIST A

LIST B

Anode

Cathode

Electrolyte

Electrolysis

Electrodes

Non- electrolyte

Strong electrolyte

Weak electrolyte

Chemical equivalent

Electrolytic reduction

Decomposition of substance in molten or in solution by passage of electric current

A reduction process by the use of carbon monoxide or hydrogen

Relative concentration of ions in solution.

Two poles of carbon or metal dipped in solution allow electrons enter or leave electrolyte

Coating metal with a thin coat of another metal by electrolysis

Positively charged electrode which electrons enter the external circuit.

Decomposition of chemical

Substance which will conduct electricity.

Hydrogen gas liberated due to the passage of electric current

Negative electrode

Extraction of most reactive metals

Oxygen gas in collected

Splits up during electrolysis

Faraday constant divide by its valency

Atomic mass divide by its valency.

Completely dissociate during electrolysis.

Partially dissociate during electrolysis.

Copper metal is deposited

Relative position in electrochemical series.

Substance that do not conduct electricity.

(a) Define

Electrochemical equivalent

Chemical equivalent

(a) State the chaterliers principle

(b) With the help of an energy level diagram on the reaction

CaCO3 ⇋ CaO(g) + CO2(g) ∆H = + 175.5 kJ. Mol -1

Show weather the reaction is exothermic or endothermic

Write down the type of the reaction shown by the above equation

(i) The temperature is increased?

(ii) The pressure is increased?

(a) Explain the meaning of the following terms;

Electrolysis

Electroplating

(b) State faraday laws of electrolysis.

electrolysis of dilute sulphuric acid using platinum electrode.

(ii) 19300 coulombs of electricity was passed through a solution of copper (II) nitrate.

Calculate the mass of copper deposited at the cathode.

(a) (i) State faradays first law of electrolysis

(ii) What is an electrolyte?

(b) What volume of oxygen measured at STP will be produced during electrolysis by a current of 5A supplied for 30 minutes?

A dilute solutions of magnesium sulphate was electrolysed using carbon electrons electrodes.

Cathode

Anode

Comment on the changes of concentration of hydration of hydrogen ions in the electrolyte.

If 3.2 of element M were deposited by an electric current of 0.5 A Passed for 90 minutes:

Calculate the quantity of electricity used.

Calculate the quantity of electricity used to deposit one mole of element M.

(M= 226, 1F = 965000C)

Work out on the charge on the ion M.

Copper electrodes are used in the electrolysis of copper (II) sulphate solution. Write down the half equations for the reactions that take place a the:

Anode

Cathode

TOPICAL QUESTIONS ON ELECTROLYSIS

FORM THREE CHEMISTRY.

SECTION A. 20 MARKS.

1.MULTIPLE CHOICE QUESTIONS

During electrolysis the molten Aluminium oxide, 3 Faradays were needed to deposit one mole of aluminium. The number electrons of aluminium will be :

6.02 x 1023

1.806 x 1023

18.06 x 10.23

180.6 x 1023

1806 x 1023

If a steady current of 2 amperes were passed through an aqueous solution of iron deposited are the cathode will be:-

54g

56g

0.54g

28g

0.52g

The use of electricity to decompose molten sodium chloride into its components elements is an example of:-

Electroplating

Galvanization

Hydrolysis

Precipitation

Electrolysis

What is the meaning of electrolysis?

Decomposition of a chemical compound by electricity

Breaking up by using electric power

Oxidation reduction process

Passing electricity through a conductor

Which is the correct property of an electrolyte

It contains free charge carriers

It does not have free ions

It can not be decomposed by an electric current

It contains mainly undissociated molecules

Which of the following is not an electrolyte

Aqueous sulphuric acid

Solid common salt

Molten calcium chloride

Alcohol in water

When dilute sodium chloride solution is electrolysed using carbon electrodes which substance is collected at the cathode

Hydrogen

Sodium

Chloride

Oxygen

What are cations?

Cathode going ions

Anode moving ions

Negatively charged ions

Unionized particles

Which particles will be anions in the electrolysis of aqueous copper sulphate using platinum electrodes?

SO42- and OH

Cu 2+ and H+

Cu 2+ and SO4 2-

Platinum

What is the name given to quantity of substance deposited by one coulomb of electricity on the electrode?

Electrochemical equivalent

Chemical equivalent

1 mole of the substance

1 G.M.V at s.t.p.

Which one of the following is the application of electrolysis

Electroplating

Isolation of elements

Purification of metals

All of the above

What is the name of the vessel in which electrolysis process is conducted

Voltameter

Voltmeter

Avometer

Beaker

Match the items in list A with those from list B by writing the letter of the correct response from List B besides the item number in List A.

LIST A

LIST B

Anode

Cathode

Electrolyte

Electrolysis

Electrodes

Non- electrolyte

Strong electrolyte

Weak electrolyte

Chemical equivalent

Electrolytic reduction

Decomposition of substance in molten or in solution by passage of electric current

A reduction process by the use of carbon monoxide or hydrogen

Relative concentration of ions in solution.

Two poles of carbon or metal dipped in solution allow electrons enter or leave electrolyte

Coating metal with a thin coat of another metal by electrolysis

Positively charged electrode which electrons enter the external circuit.

Decomposition of chemical

Substance which will conduct electricity.

Hydrogen gas liberated due to the passage of electric current

Negative electrode

Extraction of most reactive metals

Oxygen gas in collected

Splits up during electrolysis

Faraday constant divide by its valency

Atomic mass divide by its valency.

Completely dissociate during electrolysis.

Partially dissociate during electrolysis.

Copper metal is deposited

Relative position in electrochemical series.

Substance that do not conduct electricity.

(a) Define

I) Electrochemical equivalent

Chemical equivalent

(a) Explain the meaning of the following terms;

Electrolysis

Electroplating

(b) State faraday laws of electrolysis.

(c) (i) Write chemical equation for the reaction that take place at each electrode during the electrolysis of dilute sulphuric acid using platinum electrode.

(ii) 19300 coulombs of electricity was passed through a solution of copper (II) nitrate.

Calculate the mass of copper deposited at the cathode.

(a) By the help of a well labelled circuit diagram explain the electrolysis of a dilute

solution of sodium chloride using graphitic carbon electrodes.

(b) Is the above electrolysis having special name? If yes explain.

dryness of the electrolyte? Explain your answer.

Calculate the mass of magnesium metal that will be produced during the electrolysis of molten magnesium chloride if a current of 1.93A is passed through for 16 minutes and 40 seconds.

(a) State faradays 1st law of electrolysis

(b) What mass of silver and what volume of oxygen (at s.t.p) would be liberated in electrolysis by 9650 coulombs of electricity?

7. I)(a) State the faraday law’s of electrolysis

(b) An element P has relative atomic mass of 88. When current of 0.5A was passed through fused chloride of P for 32 minutes and 10 seconds, 0.44g of P was deposited at cathode;

Calculate the number of faraday required to liberate 1mole of P.

i).Write the formula of P ions

ii).Write the formula of hydroxide of P

(a) Distinguish between chemical equivalent and electrochemical equivalent of a

substance.

(b) Describe electrolysis of copper (II) sulphate using graphite electrodes.

Note: Diagram is necessary.

I) (a) State Faraday laws of electrolysis.

(b) An element Z has a relative atomic mass of 88. When a current of 0.5 amperes was passed through fused chloride of Z for 32 minutes and 10 seconds 0.44 of Z were deposited at the cathode;

Calculate the number of Faradays needed to liberate one mole of Z.

Write the formula of the Z ions

Write the formula of hydroxide of Z.

(a) i. List down three (3) factors affecting the selection of ion discharge at the electrode.

ii. Define the term electrolyte.

(b) A bluish copper sulphate aqueous solution was electrolysed by using copper

electrodes.

Write ionic chemical equations for the reactions which occurred at the cathode and anode.

Explain what will happen to blue colour of copper sulphate solution as electrolysis continues.

The following apparatus was used in experiment to electroplate an iron knife with Silver;

The K.Ag(CN)2 contains Ag+, K+ and CN ions;

ii. Which electrode is the cathode?

Name the process taking place at the anode as either oxidation or reduction.

Represent the process at each electrode by the appropriate ionic equation.

What happen to the electrolyte?

(a) 0.02 moles of electrons were passed through a solution of sodium hydroxide using platinum electrodes.

Give the names of the gases evolved to each electrode.

Write ionic equations of the reaction taking place at the electrodes.

Calculate the number of moles of each gas produced and the volume which each gas would occupy at S.T.P.

(b) What mass of copper will be liberated during electrolysis of copper sulphate solution by a charge of one faraday?

(c) An element X has a relative atomic mass of 88. When a current of 0.5A was passed through the fused chloride of X for 32 minutes 10 seconds, 0.44g of X was deposited at the cathode.

Calculate the number of faradays needed to liberate 1 mole of X.

Write the formula for X ion.

Write the formula for hydroxide of X.

SAMPLE NECTA QUESTIONS.

2004.

4. (a) Sodium, magnesium, zinc, copper and silver are five metals which appear in this order in the activity series; sodium being the most reactive and silver the least reactive. Which one of these metals is:

Likely to tarnish most rapidly when exposed to air?

Most likely to be found free in nature?

Least likely to react with steam?

Two of the metals in 4(a) above are usually extracted by electrolysis of their molten chlorides. Name the two (2) metals and give one reason of using this method.

(1) Name the positive and negative electrodes of an electrolytic cell.

To which electrode will sodium ions in an aqueous dilute solution of sodium chloride migrate during electrolysis?

What other ions will migrate to the electrode stated in 4(c)(ii)?

Which ions will be discharged at the electrodes stated in 4(c) (ii)?

Give reasons for your answer.

2005.

6. (a) (i) List down three (3) factors affecting the selection of ion discharged at the electrodes.

(ii) Define the term electrolyte.

A bluish copper sulphate aqueous solution was electrolyzed by us ::-_i copper electrodes.

0) Write ionic chemical equations for the reactions which occurred at the cathode and anode.

(ii) Explain what will happen to the blue colour of copper sulphate II electrolysis continues.

How many moles of electrons are required to produce 27g of Al the electrolysis of molten Al203?

2006 QN 7.

State Faradays:

First law of electrolysis.

Second law of electrolysis.

(b) Explain the meaning of:

Electrolysis

Electroplating

c) (i) Write a chemical equation for the discharging process at the anode and cathode when dilute sulphuric acid is electrolyzed using platinum electrodes.

(ii) 0.2 Faraday of electricity was passed through a solution of copper sulphate. Calculate mass of copper deposited

2013.

QNS 8. (a) A current of 0.5A was made to flow through silver voltmeter for 30 minutes. Calculate the mass of silver deposited and the equivalent weight of silver.

2013.

QNS 12 Assume that you are a chemist in a chemical plant that deals with production of chlorine gas. You want to produce 100 litres of chlorine gas per hour so that you can reach the companys goal of producing 2400 litres every day. What current of electricity will you allow to flow per hour?

2016.

QNS6. (a) Write the formula for the following compounds:

(i) But-2-ene

(ii) Pent-2-yne

(iii) 1, 2 dichloroethane

(iv) 2, 4 dimethylhexane

(b) Briefly explain what will be observed when silver nitrate solution is added to aqueous solution of sodium chloride.

2016.

QN 13. 0.48 of a metal, M was placed in a test tube and hot copper (II) sulphate solution was added to it and stirred until the reaction stopped. The metal (M) displaced copper from copper (II) sulphate solution. Copper was filtered washed with water, dried at 100°C and the mass found to be 1.27g. Given that the balanced chemical reaction that occurred is M + CUSO4(aq) → MSO4(aq) + Cu(s)

(a) Calculate:

The number of moles of copper that were formed and the number of moles of M that were used in the reaction.

The relative atomic mass of M and hence identify metal M.

(b) State the appearance of the metal formed (Cu).

TOPIC : 8 COMPOUNDS OF METALS

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

COMPOUNDS OF METALS.

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

SECTION A. 20 MARKS.

MULTIPLE CHOICE QUESTIONS

Which of the following properties generally increases down the group?

Ionization energy

Atomic sixe

Electro negativity

Sodium and zinc

When does a chemist fail to identify a compound between sulphur and

iron?

a black solid is formed

heat is used to join them up

yellow color of sulphur and silvery shinny

the resulting mass is greater than the individual mass of the elements

any of the above A-D does not take place.

Which of the following sets of elements is arranged in order of increasing electro negativity.

A. Chlorine, fluorine, nitrogen, oxygen, carbon

B. Fluorine, chlorine, oxygen, nitrogen, carbon

C. Carbon, nitrogen, oxygen, chlorine, fluorine

D. Nitrogen, oxygen, carbon, fluorine, chlorine

E. Fluorine, nitrogen, oxygen, chlorine, carbon

(iv) Which method could be used to separate the products in the following equation?

Pb(NO33)2(aq) + 2KI(aq) → Pb12(s) + 2KNO3(aq)

Colourless colourless yellow colourless

A. Chromatography B. Crystallisation

C. Distillation D. Filtration E. Condensation

(v) The metal nitrate which will NOT give a metal oxide on heating is

calcium nitrate

silver nitrate

lead nitrate

copper nitrate

zinc nitrate

(vi) Which of the following pairs of compounds can be used in preparation of calcium sulphate?

Calcium carbonate and sodium sulphate

Calcium chloride and ammonium sulphate

Calcium hydroxide and barium sulphate

Calcium nitrate and lead (II) sulphate

Calcium chloride and barium sulphate

(vii) The ionic equation when aqueous ammonium chloride reacts with sodium hydroxide solution is represented as:

NH4(aq) +0H-(aq) +0H (aq) → NH3(g) + H20(I)

Na+(aq) + Cl-(aq) → NaCl (s)

H+(aq) + OH(aq) → H2O(l)

2NH4(aq) + 2Cl-(aq) → 2NH3(g) + 2HCl(g)

(viii) Three elements, X, Y and Z, are in the same period of the periodic table. The oxide of X is amphoteric, the oxide of Y is basic and the oxide of Z is acidic. Which of the following shows the elements arranged in order of increasing atomic number?

X, Y, Z

Y, Z, Y

Z, X, Y

Y, X, Z

X, Z, Y

(ix) The only metal which does not react with dilute hydrochloric acid is:-

Magnesium

Aluminium

Copper

Zinc

Sodium

(x) Brown ring test is a technique used to test the presence of;

A. Sulphate

B. Nitrate

C. Chloride

D. Potassium salts

Match the items in List A with their correspondence responses in List B.

LIST A

LIST B

It nitrate is used in manufacture of antifungal cream.

Its carbonate is used in removing water hardness.

Its nitrates is used as fertilizers.

Its hydroxide is used to make ant - acid tablets.

Brown ring test.

Its nitrates form black residue when heated.

Its hydroxide is used as fuming agent in fire extinguisher.

Basic oxide

Acidic oxide

Yellowish solid produced when sodium burn in excess air.

Ammonium radical

Sulphur

Carbon

Copper

Sodium oxide

Sodium peroxide

Silver

Sodium

Gold

Aluminium

Phosphorous pent oxide

Aluminum oxide

Dinitrogen monoxide

Magnesium

Germanium

Iron (ii) sulphate and sulphuric acid.

Argon

Potassium zincate

Test of carbonate in the laboratory

SECTION B

A solution of chlorine in water is acidic.

Yellow phosphorus is stored under water.

Give an account of the following:

Anhydrous copper (II) sulphate becomes coloured when exposed to the air for a long time.

Carbon dioxide can be collected by the downward delivery method

Concentrated sulphuric acid is not used for drying hydrogen sulphide gas.

Sodium metal is kept in paraffin oil.

4. (a) With the aid of chemical equations, explain what will happen when aluminum chloride reacts with water.

(b) A student accidentally broke a beaker containing copper (II) sulphate crystals. He decided to separate the blue crystals from the small pieces of glass by first dissolving the mixture and then filtering. What were the next steps?

5. (a) Asubuhi Njema child was sick. When she took her to the hospital, she was prescribed some medicine including a bottle of syrup. The bottle was written: shake before you use. What does this statement signify?

(i) What is the first step to take when you want to identify the contents of a given salt containing one anion and one cation?

(ii) In a solution of salt and water, identify a solute and a solvent Justify your answer.

Sodium is a solid while chlorine is a gas at room temperature although they are in the same period in the Periodic Table. What is the cause of this difference?

6. (a) Describe the effect of:

Strongly heating a piece of marble in a Bunsen burner flame.

Moistening the residue from (i) above with water.

(i) For what reason is slaked lime added to the soil in gardening?

(ii) Why is concentrated sulphuric acid used as a drying agent?

Why is zinc used as a coat for iron and not vice versa?

(d) What will happen when:

Yellow flowers are introduced into a gas containing chlorine gas?

A burning splint is introduced into a gas jar containing hydrogen gas?

A glass rod which was dipped in concentrated hydrochloric acid is introduced into a gas jar containing ammonia gas?

Sulphur dioxide gas is bubbled through a yellow acidified potassium dichromate solution?

7. (a) Name the ore commonly used in the extraction of copper metal.

(b) Steps (i) to (iv) below are used during the extraction of copper metal from its ore. Write a balanced chemical equation for each step.

(1) Roasting of the concentrated ore (CuFeS2 ) in air.

Heating the roasted ore with silica in the absence of air.

Burning copper sulphide in regulated supply of air.

Purification of copper by electrolysis using copper sulphate solution electrolyte, pure copper cathode and impure copper obtained from the extraction anode.

An iron earring dropped into a container of copper sulphatesolution.

Copper knife dipped into zinc nitrate solution.

Copper turnings dropped into a container of dilute hydrochloric acid.

8. (a) Explain briefly the following observations with the help of equations.

i) White anhydrous copper (II) sulphate changes its colour to blue when water is added.

Vigorous reaction takes place when a small piece of sodium metal is placed in water.

Addition of Zinc metal into a solution of copper (II) sulphate results into decolourization of the solution and deposition of a brown solid substance.

9. (a) Sodium, magnesium, zinc, copper and silver are five metals which appear in this order in the activity series; sodium being the most reactive and silver the least reactive. Which one of these metals is:

Likely to tarnish most rapidly when exposed to air?

Most likely to be found free in nature?

Least likely to react with steam?

Two of the metals in 4(a) above are usually extracted by electrolysis of their molten chlorides. Name the two (2) metals and give one reason of using this method.

(1) Name the positive and negative electrodes of an electrolytic cell.

To which electrode will sodium ions in an aqueous dilute solution of sodium chloride migrate during electrolysis?

What other ions will migrate to the electrode stated in 4(c)(ii)?

Which ions will be discharged at the electrodes stated in 4(c) (ii)?

10. (a) Chemical reactions can be classified into several types. Study the following reactions and then indicate to which type they belong.

i) Action of heat on lead nitrate

ii) The addition of dilute hydrochloric on silver nitrate

Action of heat on calcium carbonate in a closed vessel.

iv) Action of chlorine gas on iron(II) chloride

Explain how you can separate a mixture of iodine and copper (II) oxide.

List three product formed when copper (II) nitrate is heated strongly in a test tube.

Aluminum has atomic number 13.

Write down the formula of aluminum chloride.

State and explain the observation made when a burning magnesium ribbon is plunged into a gas jar of carbon (IV) oxide.

State why water is not a suitable extinguisher of an e the observation made when a burning magnesium ribbon is plunged into a gas jar of carbon (IV) oxide.

State why water is not a suitable extinguisher of an electrical fire.

State the name given to the reaction in which ethane is converted to ethane.

Calculate the volume of 0.09M acid that can be from 1DM3 of 1.2M acid.

Explain how you would distinguish between a sulphate using barium choride solutions and silute hydrochloride acid.

15.2g of the oxide of metal X was formed when 10.64g of X reacted with excess oxygen. Dermine the formula of the metal oxide of X . (X = 56, O = 16)

Name the products found at the cathode anode when concentrated sodium chloride is electrolytes using carbon electrodes.

Some iodine became contained with black copper (II) oxide.

Explain how you would separate the mixture.

Name the chemical techniques that can be used to separate sodium chloride is electrolyzed using carbon electrodes.

Name a substance which decomposes on heating to give two gaseous, other than water vapor.

Name the two gaseous above.

A compound J React with concentrated sulphuric acid (H2SO4) TO GIVE BROWN fumes of a compound K. a solution K in water reacts with copper metal to give a dark brown gas L. Name substances J, K and L.

School Base-OnlinePage 6

TOPIC : 8 EXTRACTION OF METALS

CHAPTER 8

EXTRACTION OF METALS

KEY TERMS AND CONCEPTS.

Metal- this is an element that reacts by loosing electrons

Metallurgy- this is the study of properties, production and purification of metals

ore- this is a mineral that contains a large proportions of metal or metallic compound that has an economic value.

Dressing- this is removing of impurities from ores of metal without decomposing it.

Calcinations- is the process in which the ore is heated in the absence of air below its melting point to expel water from hydrated oxide.

Roasting- is a process in which the ore is heated in the presence of air at temperature below the melting point of the ore.

Blast furnace- this is an apparatus that is used to extract iron from its ore.

Slag- a mixture of silicon dioxide and other impurities formed in the blast furnace and prevents the oxidation of iron formed.

Pig iron- impure variety of iron ore obtained from the blast furnace.

Cast iron- moulds of desired shaped formed from molten iron which has been re-smelted.

Chemical strength- this is the measure of the reactivity of a metal, which is depended on its ability to loose electrons.

Tensile strength- this is the ability of a metal to withstand force applied on it without breaking.

Reactivity series- this is arrangement of metals according to their reducing power.

Displacement reaction- this is a type of reaction in which a less reactive element is replaced by a more reactive element from a compound.

Down cell- the apparatus that is used in the extraction of sodium from its ore

Brine- this is concentrated sodium chloride.

Introduction

Metals are very useful. Metals are obtained from ores.

Ores are naturally occurring rocks that contain metal compounds. They have a wide range of applications in industries and at home.

Ores are only considered worthy for mining if they contain sufficient amount of metal compounds.

Ores are only considered worthy for mining if they contain sufficient amount of metals compounds.

The method used to extract metals from the ore in which they are found depends on their reactivity and physical properties.

For example reactive metals such as sodium are extracted by electrolysis while less reactive metals such as iron may be extracted by electrolysis while less reactive metals such as iron may be extracted reduction with carbon (II) Oxide.

The physical properties of metals considered during the separation from their ores include; Density, Particle size and shape, electrical and magnetic properties.

Occurrence and location of metals ores in Tanzania

Tanzania is endowed with many mineral resources. She has a great potential particularly for gold, copper, cobalt and various ferrous minerals.

The table below shows different metals and their distribution within Tanzania.

Distribution of metal Ores in Tanzania

Site Name	Province	Metal	Development Status
Kikabwe	Dodoma	Manganese	Occurrence
Hundus Deposit	Morogoro	Iron Titanium Vanadium	Occurrence
Liganga Deposit	Iringa	Iron Titanium Vanadium	Occurrence
Almasi Mine	Shinyanga	Diamond	Producer
Nkenze Chromium Deposit	Iringa	Chromium	Occurrence
Chambogo	Kilimanjaro	Magnesite	Producer
Kabanga		Copper Cobalt Nickel	Prospect
Liganga	Iringa	Vanadium	Occurrence
Buck Reef	Mwanza	Gold Molybdenum	Producer
Hiso- Haneti	Dodoma	Silica	Producer
Bunduki	Morogoro	Manganese	Occurrence
Mombo	Tanga	Aluminum	Occurrence
Uruwira	Rukwa	Silver Gold	Producer
Kigugwe Copper Prospect	Mbeya	Copper	Prospect

Extraction of metals

The branch of science concerned with the properties, production and purification of metals is called Metallugry.

Metals are extracted from their ores. An Ore is a naturally occurring solid material from which a metal or valuable mineral can be extracted.

The main processes involved in the extraction of metals from their ores are:

Ore purification

Extraction of the Metal

Refining of the Metal.

Ore purification

Ore purification is done in order to remove the impurities present in the metal ore. The methods used in purifying metal are ores include:

Dressing

Calcinations

Roasting.

Dressing

Dressing is the removal of impurities from the metal ore without decomposing the ore chemically.

The main impurities which are removed during the dressing of metal ores are sand, limestone, quartz and silicate. Dressing is also known as concentration.

Calcination

Calcination is a process in which the ore is heated in the absence of air below its melting point to expel water from a hydrated oxide or expel carbon dioxide from a carbonate.

In calcinations, small molecules such as water, carbon dioxide or sulphur dioxide are usually expelled from the ore.

Roasting

Roasting describes a process in which the ore is heated in the presence of air, usually at temperatures below the melting point of the ore.

The ore may be heated alone or mixed with other substances.

Roasting usually involves greater chemical changes such as oxidation or chlorination. Both calcinations and roasting take place in the furnace.

Extraction

The methods used in extracting metals from their ores include reduction by heating in the air (autoreduction), reduction by carbon (smelting), reduction by carbon (smelting) reduction by electrolysis, reduction by precipitation and amalgamation.

The choice of the reducing metal used for a given metal or depends on the position of the metal ore depends on the electrochemical series. The table below shows the methods used in reducing different metal ores.

Method used to reduce different metal ores

Method	Metal/ metal ore
Reduction using carbon smelting	Oxidation of less electropositive metals such as lead, zinc, iron, tin and copper.
Reduction by heating in air (autoreduction)	Oxides and sulphides of less electropositive metals such as mercury copper and lead. Cinnabar (HgS) is reduced to mercury using this method.
Reduction by electrolysis	Oxides, hydroxides or chlorides of metals high in the reactivity series such as Sodium, Calcium, Magnesium, and Aluminum.
Reduction by precipitation	Used to extract less reactive metals a solution containing the metal ions is reacted with a more reactive metal displaces the less reactive metal from the solution.
amalgamation	Used to extract less reactive metals like silver and gold. The finely crushed ore is brought into contact with mercury which forms an alloy (amalgam) with the metal

Table 8.1

Refining or purification

The metals obtained through the reduction process are generally impure. They are known as crude metals.

Crude metals contain impurities such metals. Crude metals contain impurities as the other metals, non metals such as silicon and phosphorus and unreduced oxides and sulphides of metals.

Sometimes impurities are also introduced in the process of when preparing the metal ore from reduction.

This is especially the case when the chemical method is used to concentrate theory. The three main methods used to purify the crude metal are distillation, oxidation and electro- refining.

Distillation

The crude metal is heated in a furnance until the pure metal evaporates, leaving behind the impurities.

The vapour is then collected and condensed in a separated chamber.

Oxidation

Oxidation is the addition of oxygen. The molten crude metal is exposed to air in a furnace.

The impurities are oxidized and escape as vapour or from a scum over the molten metal.

The scum is then removed by skimming.

This method is used only when the impurities have a great affinity for oxygen than the metal.

Electro- refining

Most metals are used purified by electrolysis (electro- refining).

The crude metal is molded into blocks and made the anode of an electrolytic cell. The cathode is usually made up of a thin plate of on the cathode.

The soluble impurities go into solution while the insoluble impurities settle down below the insoluble impurities settle down below the anode as anode mud or sludge.

Extraction of Sodium

Occurance

Sodium does not occur naturally as a free element. This is because it is too reactive. However, it is abundant in different compounds such as sodium chloride (rock salt) and sodium sulphate in sea water.

Extraction

Sodium is extracted by the electrolytic reduction of purified molten rock salt (sodium chloride).

The electrolysis is mainly carried out in the Down’s cell shown in figure 6.2. The cell has a steel cathode and a graphite anode.

Calcium chloride is added to sodium chloride to lower the malting point of sodium chloride from 7740C to about 6000C.

The composition of the electrolyte is 40% sodium chloride and 60% calcium chloride.

Chloride ions move to the anode while sodium ions move to the cathode.

Figure 8.1

The reactions taking place at the electrodes are shown below:

Anode: 2Cl-

Cl2 (g) + 2e-

Cathode: Na+ + e-

Na (I) Na (I)

Chloride gas is therefore formed at the anode while sodium metal is produced at the cathode.

At 6000C, the sodium and chloride produced could react violently if allowed to come into contact. To prevent this cell has steel gauze (diaphragm) around the anode to keep the two products apart.

The large graphite anode is used to facilitate maximum oxidation of chloride ions to chlorine gas.

This also maximizes the formation of sodium metal at the cathode. Chloride collects in the inverted trough (hood) placed over the cathode, rises up the pipe and is tapped off through an opening on the upper part of the Downs cell because of its low density which makes it to float over the mixture.

The sodium metal is collected from upwards in the Downs cell because of its low density which makes it to float over the mixture.

The sodium metal from the Downs cell contains some calcium, which is also formed through electrolysis. The calcium crystallizes when the mixture cools and a relatively pure sodium metal is obtained.

Extraction of Aluminum

Occurance

Aluminium is the most abundant metal on the earth crust , making about 8% of the crust. Aluminum does not occur naturally as a free element.

It occurs in the form of oxides, fluorides, silicates and sulpahates. The following are the main ores of aluminum:

Bauxite (Al2O3. 2H2O)

Cryolite (Na3AlF6).

Felspar (KAlSi3O8)

Kaolin (AL2Si2 O7. 2H2O)

Bauxite and cryolite are the chief ores from which alluminium is extracted.

Extraction

Aluminium is extracted from bauxite by electrolysis. This is done using the hall heoroult method. The following are the stages involved:

Purification of bauxite.

Electrolytic reduction of alumina(Al2O3)

Refining of aluminium .

Purification of bauxite

Bauxite contains oxides of iron and silicon as impurities. These impurities are removed using the Hall process where the powdered ore is heated to bright redness with oxide is amphoteric, it dissolves to form sodium aluminate (Na AlO2):

Al2O3.2H2O (s) + Na2CO(s)

2NaAlO2(I) + 2H2O(I) + CO2(g)

The insoluble iron and silicon oxides are left as residue.

The molten mass is removed with the water while the insoluble iron and silicon oxides are left behind as residue. The filtrate is heated to between 500 and 6000c. A stream of carbon dioxide is then passed through the mixture, resulting in the precipitation of aluminum hydroxide.

2NaAlO2 (I) +3H2O(I) + CO2(g)

2Al(OH)3 (s) + Na2CO3(aq)

Figure 8.2

Aluminium Hydroxide is filtered off, washed, dried then calcinated at 15000C to give alumina.

2AL (OH)3

Al2O3 (I) + 3H2O (I)

Electrolysis of alumina (Al2O3)

Alumina cannot be reduced to aluminum metal by Carbon because aluminum has a greater affinity for oxygen than carbon.

Alumina also reacts with carbon at high temperatures to form Alumina carbonate (Al4 C3).

The electrolysis of alumina is carried out in the presence of cryolite (Na3 AlF6). This improves the electrical conductivity of alumina; a which is otherwise a bad conductor of electricity.

It also lowers the melting point of alumina, which is otherwise a bad conductor of electricity.

It also lowers the melting point of alumina. If the electrolysis is carried out at very high temperatures, the alumina formed would vaporize.

The cell consists of an iron bath lined with graphic which act as a carbon which act as the cathode.

The anode consists of number of carbon rods attached copper clamps and suspended in the fused mass.

The carbon roads are arranged in such a way that they can be raised, lowered or replaced when there is a need.

The reactions taking place at the electrodes are shown below:

Cathode: Al3+ (I) +3e-

Al (I)

Anode: 2O2-(I)

O2 (g) + 4e_

Once discharged, the Aluminium collects at the electrolysis is an expensive process. The temperature of the fused electrolyte has to be maintained at 8000C – 9000C.

This is achieved by passing a steady current of about 100 000 A through the electrolyte.

The process is only economical where electricity is cheap and readily available.

Purification of crude aluminium

The aluminium obtained from the Hall – heroult process is 99% pure. It contains small quantities of iron , silcon, alumina and carbon.

It is purified futher by electrolysis using Hoopes electrolytic cell.

The Hoope’s electrolytic cell is made up of an iron tank lined on the inside with carbon which serves as the inside with carbon which serves as the anode. The tank has three different layer s of molten substance:

The bottom layer contains molten impure aluminium containing copper and slicon to increase the density.

The middle layer contains a mixture of fluorides of sodium, Barium and aluminium in fused form. This layer acts as the electrolyte. The top layer contains pure molten aluminum which together the carbon rods serve as the cathode. The carbon cathodes are suspended from above.

HOOPE’S ELECTROLYTIC CELL DIAGRAM

Figure 8.3

On Passing an electric current the aluminium formed in the middle layer passes top layer as pure aluminium passesfrom the bottom layer to the middle layer and the process continues.

Aluminium of about 99.99% purity is tapped from the tapping hole.

Extraction of Iron

Iron is the second most abundant metal occurring on the earth’s crust after aluminium. It does not occur as a free element because it is relatively reactive.

The following are the most common cores of the iron

Haematite (Fe2O3)

Magnetite( Fe3O4)

Limonite (2Fe2O3. H2O)

Siderite or spathic iron ore (FeCO3)

Iron pyrites (FeS2)

The most important ores of iron are hematite (Fe3O4) which contains 70 % iron and magnetite (Fe3O4) which contains 72.4% iron. Iron pyrite (FeS2), although abundant on the earth’s crust is not used as a source of iron. It is mainly used in the production of sulphuric acid.

Iron is mainly extacted from the haematite ore by reduction using carbon (coke). The following are the stages involved in the extraction process:

Concentration

Roasting

Smelting (reduction)

Concentration

The haematite ore is first dressed to remove earthy impurities such as clay and sand , and non magnetic impurities.

Roasting

The concentrated ore is heated strongly in excess air in shallow kilns. The following changes occur during the roasting process:

Most of the moister is removed.

Fe2O3. xH2O(g)

Fe2O3(S) + XH2O(I)

Sulphur and arsenic impurities are oxidized to their oxides. They escape as gases.

S(S) + O2(g)

SO2 (g)

4As(s) + 3O2

2As2O3(g)

Carbonates are decomposed into oxides giving off carbon dioxide gas. For example,

FeCO (s)

FeO(s) + CO2(g)

Iron (II) Oxide Is Oxidized To Iron (III) oxide.

4FeO(s)+ O2(g)

2Fe2O3(S)

This reduces the loss of iron since iron (II) oxide forms a slag with silicon dioxide forms a slag with silicon dioxide (sand)

FeO(s) + SiO2(S)

FeSiO3(s)

The entire mass becomes porous and suitable for reduction to metallic iron in the blast furnace.

Smelting

The roasted ore is smelted, that is reduced by carbon in the blast furnance. The blast furnace is a tall structure, about 30 m high and 10 m in diameter at its widest part.

The furnace contains a firebrick lining inside a steel shield. The ore is mixed with carbon (coke) and limestone (CaCO3) and introduced into the furnace from the top. Blasts of hot air are blown into the furnace through small opening near the base known as tuyeres.

DIAGRAM FOR BLAST FURNACE

Figure 8.4

The blast furnace is divided into three zones, according to the temperatures in the furnace.

The upper zone of reduction

This zone is near the top of the furnace. The temperature range in this zone is 3000C – 8000C. In this zone the carbon dioxide.

C(s) + O2(g)

CO2(g)

The carbon dioxide formed reacts with more coke to form carbon monoxiede:

CO2 (g) + C(s)

O(g)

The carbon monoxide formed reduces iron(III) Oxideto spongy iron.

Fe2 CO3 (s) + 3CO(g)

2Fe(I) + 3CO(g)

Note that the main reducing agent in the blast furnace is carbon monoxide and not carbon.

Since the temperature in this zone is too low melting iron, yhe metal obtained is called spongy iron.

As the iron ore is being reduced, part of the limestone decomposes at 6000C

to lime(CaO):

CaCO3(S) + C(s)

CaO(s) + Co(g)

The lime stone also reacts with coke to form lime.

CaCO3(s) + C(s)

CaO(s) + 2CO (g)

The middle zone

This is the lower zone of reduction. The temperature range in the zone is 9000C-1200OC. The following are reactions that take place in this Zone.

Carbon dioxide is reduced to carbon monoxide by coke:

CO2 (g) + C(s)

2CO(g)

This reaction is accompanied by the heat absorption

At about 10000C, Calcium carbonate is almost completely decomposed to lime (CaO):

CaCO3 (s)

CaO(s) + CO2(g).

The calcium oxide act as the and combine with the silica (SiO2)(s)

CaSiO3(I)

A flux is a mixture of chemicals that react with impurities to form Slag.

The lower zone of combustion

This is the lowest and the hottest part of the furnace. The temperature rang is 12000C- 15000C. The following reactions take place in this zone.

Coke burn in the blast of hot air to produce carbon dioxide:

C(s) + O2 (g)

CO2(g)

The spongy iron melts in the zone at about 13000C and collects at the bottom of the heart.

Iron (II) oxide which might have escaped the reduction process in the middle zone is reduced by coke in this zone.

2FeO(s) + C(s)

2Fe (s) + CO(g)

The bottom of the furnace, the molten iron sinks while the fusible slag, being less dense, floats over the floated iron forming being oxidized by the hot air. The slag and iron are periodically removed through different outlets.

The mixture of waste gases containing nitrogen, carbon dioxide and carbon monoxide are known as blast furnace gases. The mixture is burnt in air to produce heat which is used for heat in pre- heating the blast coming in through the tuyeres.

The iron obtained from the blast furnace is an impure variety known as pig iron. Pig iron is further purified by re- smelting it with coke and lime in another furnace called cupola. The molten iron from the cupola is poured into moulds of desired shape. The iron thus obtained is called cats iron.

Chemical properties of metals

All metals behave in a similar way when chemically combined with other substances. In chemical reaction, metals give electrons to other elements.

The tendency to lose electrons and form positive ions is called electro positivity. In solutions, metals form positive ions (cations).

Physical Strength

The physical strength of a metal is the ability of the metal to withstand force applied on it without breaking. The physical strength of a metal is referred to as its tensile strength. Iron has a very high tensile strength while sodium has a low tensile strength.

Wires of the same gauge, but made from different metals, support different loads (masses) before changing from being elastic to being plastic.

In the elastic state, metals undergo elastic deformation when stress is applied. Elastic deformation is recoverable after the load is removed. Plastic deformation is not recoverable.

The point at which material changes from elastic deformation to plastic deformation is a measure of its tensile strength this point is referred to as the yield strength off the material.

Chemicals strength

The chemical strength of the metal is measure of how readily the metal takes part in chemical reaction.

Some metals are very reactive, readily taking part in chemical reaction to produce new substances. Potassium and sodium are some of the most reactive metals.

Metals such as platinum are very uncreative. Such metals do not take part in chemical reaction easily.

Calcium is quite reactive with cold water forming calcium hydroxide, which is moderately soluble , and hydrogen, gas.

Magnesium reacts slowly with cold water forming magnesium hydroxide, which is slightly soluble in water, and hydrogen gas. The reaction is faster with boiling water. The reaction is faster with boiling water.

Aluminum dies not react with water, even when heated in steam. This is because of the formation of an oxide layer (Al2 O3) which prevents further reaction with the metal.

Iron does not react with both cold and boiling water. However, when the metal is heated in steam, an oxide of iron and hydrogen gas are formed.

Copper does not react with water, not even steam.

Redox reactions involving metals

In Redox reactions, oxidation and reduction take place simultaneously.

Reactions involving oxidation and reduction are known as redox reactions.

Oxidation is a process which involves the loss of one or more electrons from a substance. For example, sodium metal loses an electron to form a sodium iron:

Na(s)

Na+ + e-

In above reaction, sodium has been oxidized to sodium ion. An equation which shows the loss of electrons is referred to as an oxidation half reaction.

Reduction, on the other hand involves the gaining of or more electrons. for example, chlorine gas can be reduced to chloride ions through gaining electrons:

Cl2 (g) + 2e-

Cl-

A redox reaction is made up of two or half reactions, one representing oxidation and the other representing reduction.

Consider the reaction between sodium metal and chloride gas to form

Sodium chloride.

2Na(s) + Cl2-

2NaCl(S)

This is a redox reaction. The two half reactions are:

2Na(s)

2Na+ +2e- (oxidation half reaction)

Cl2(g) + 2e-

2Cl-( reduction half reaction)

The reactions lost by sodium are gained by chlorine.

The substance that loses electrons is said to be a reducing agent while the substance that gains electrons is an oxidizing agent.

In this above reaction, sodium is a reducing agent while chlorine is an oxidizing.

The reducing power of the metal

In chemical reactions, metals tend to lose the electrons in the outermost energy level. As such metals are reducing agents.

The ease with metals loses the electrons increases down the group. Potassium with an electronic configuration of 2.8.8.1 is a stronger reducing agent than lithium which has an electronic configuration of 2.1.

Due to the increase in energy levels down the group the electrons in the outermost energy level get futher away from the nucleus down the group. This results in a weaker attraction to reduce the nucleus.

The electrons in the outermost energy level are therefore easily lost during chemical reactions. The reducing power of metals, thus increase down the group.

Across the period from left to right, the number of electrons in the outer most shell increases. The number of protons also increases because the number of protons is equal to the number of electrons in a neutral atom.

However, the number of shells does not change. This results in stronger attraction of the electrons in the outermost shell to the protons in the nucleus.

This makes it difficult for the electrons to be lost during chemical reactions. The reducing power of metals therefore decreases across the period from left to right. For example, sodium (2.8.1) is a stronger reducing agent than Aluminium (2.8.3).

The figure below shows the trend in the reducing power of metals down the group and across the periodic table.

Figure 8.5

Reactivity series of metals

The reducing power of metals determines its reactivity. The higher the reducing power the more reactive the metals.

The above table shows a reactivity series metals. At the top of the reactivity series is the metal with the highest reducing series is the metal with the highest reducing power (most reactive).

TABLE SHOWING REACTIVITY SERIES OF METALS

Figure 8.6

Displacement reactions takes place when an atom or group of atoms take the place of another in compound.

Consider the reaction between iron metal and hydrochloric acid to form iron ( III) Chloride.

2Fe(s) + 6HCl (aq)

2FeCl3(aq) + H2(g)

In this reaction, iron takes the place of hydrogen in the hydrochloric acid. We say that iron has displaced hydrogen from the hydrochloric acid.

A more reactive metal form a solution of one of its compounds. It is therefore possible to predict whether a metal will displace another metal from its solution by comparing their positions in the reactivity series. For example sodium will displace copper from (II) Sulphate solution.

When zinc is added to copper (II) Sulphate solution and stirred, a brown solid is formed. The beaker becomes warm indicating that the reaction is exothermic.

The brown solid is copper metal.

Cu2+ (aq) + 2e-

Cu (s)

Zinc metal dissolves to form zinc ions (Zn2+) which displace copper (II) ions from the copper (II) Sulphate solution.

In this reaction, zinc metal acts as a reducing agent, reducing copper (II) ions to copper metal. This reaction is possible because zinc is a stronger reducing agent and oxidizes zinc metal to zinc metal to zinc ions:

Zn(s)

Zn2+ (aq) + 2e-

If the set- up is left overnight, the blue color caused by copper ions fades. If the zinc powder is excess, the solution will decolorize completely.

Some displacement reactions involving metals and ions are shown in the following table.

Some of the displacement reactions involve observable color changes. The reaction between iron and copper (II) sulphate is a good example:

e(s) + CuSO4 (aq) FeSO4 (aq) + Cu(s)

The blue color of the copper (II) sulphate changes to pale green, the Colour (II) sulphate solution.

TABLE FOR REACTIONS INVOLVING METALS

Table 8.2

Carbon and hydrogen in the reactivity series

Carbon and Hydrogen although are metals included in the reactivity series. This is because they are important reference points in the extraction of metals and in the reactions of metals with water and acids.

Carbon and hydrogen are reducing agents just like metals.

Carbon

Based on its reducing power, carbon is placed between aluminium and zinc in the reactivity series. Metals below zinc can be extracted by reducing their oxides using carbon or carbon monoxide.

Metals above carbon in the reactivity series are usually extracted by electrolysis.

When copper oxide is heated with charcoal brown specks of copper metal are observed in the reaction mixture. With lead oxide, pale white speaks of lead are observed.

Charcoal (carbon) is reducing agent and reduces the oxides of copper and lead to their respective metals:

2CuO(s) + C(s)

2Cu(s) + CO2 (g)

2PbO(s) + C(s)

2Pb(s) + CO2 (g)

Carbon ( Coke) is used in the extraction of metals low in reactivity series such as copper and iron.

Hydrogen

Hydrogen is placed between lead and copper in the reactivity series. Metals below hydrogen cannot displace hydrogen from a non – metal anion.

These metals are therefore used in jewellery and in making coins. They include copper, silver, gold and platinum.

Zinc and magnesium react with hydrochloric acid give a salt solution and hydrogen gas:

Zn(s) + 2HCl (aq)

ZnCl2 (aq) + H2 (g)

Mg (s) + 2Hcl (aq)

MgCl2 (aq) + H2 (g)

The reaction between magnesium with hydrochloric acid is more vigorous than the one between zinc and the acid. This is because magnesium is higher in the reactivity series than zinc.

Copper does not readily react with dilute hydrochloric acid.

The reactions between hydrochloric acid and metals such as magnesium and zinc to give a salt solution and hydrogen gas are redox reactions.

The hydrogen in the hydrochloric acid is the oxidizing agent. It is itself reduced to hydrogen gas on these reactions.

The hydrogen in the hydrochloric acid is the oxidizing agent . It itself reduced to hydrogen gas in these reactions.

The table below shows the order of reactivity of metals with dilute acids.

Metal	Reaction
Potassium Sodium Lithium	These metals are very reactive and hence too dangerous to add to dilute acids.
Calcium Magnesium Aluminium Zinc Iron Lead	The metals react to produce a salt and hydrogen gas. Reactivity decreases down the series.
Copper Silver Gold	These metals are less reactive than hydrogen and hence do not react with dilute acids.

Table 8.3.

END TOPIC QUESTIONS

………………….. involves heating of one in the absence of air;

Dressing

Calcination

Sedmentation

Roasting

Blowing

(a) (i) Name four important ores of iron.

(ii) What objective is achieved when concentrated ore iron is roasted?

(b) Briefly describe balanced chemical reactions taking place in the blast furnace for the

extraction of iron.

Give the names of chief ores from which the following metals are obtained .

Sodium

Iron

Identify any five destructed effects of metal extraction process and suggest the appropriate measures to rectify each of the effect you have mentioned.

The figure below is downs cell used in the extraction of sodium study it and answer the questions that follow.

Identify gas x

Explain why sodium chloride is mixed with calcium chloride in the process.

What is the function of the diaghram?

Write the ionic of the occur at the:

Cathode

Anode.

State the causes of air pollution in the extraction of metals.

The figure below shows the blast furnace, study it and answer the questions that follow.

Name the substances A, B and C

What is the function of the calcium carbonate in the process?

Name the materials fed into the blast furnace from the top.

Explain the impacts of the mining on the landscape.

State 3 metals which are produced by electrolysis of molten compound of their metals.

Explain why iron products must covered with paint or other protective layer.

Most pig iron obtained from the blast furnace is converted to steel explain:

(a) Draw a well labeled diagram of Downs cell and show the half reactions in the extraction of sodium

(b) What is the effect of temperatures above 800°C in process (a) above?

Write two balanced chemical equations of the blast furnace in each of the following processes

Preparation of reducing agents

_________________________________________

Reducing of hematite

_______________________________________

Formation of slag

(i) ___________________________________________

(ii) __________________________________________

(a) Write down two uses of the slag from a blast furnace

(b) By the help three well balanced equations state the importance of preliminary roasting step in the blask furnace

(i) Importance

(ii) Equations

Write short notes about extraction of sodium in the following steps

(a) Occurrence

(b) Preparation of electrolyte

(d) Oxidation at the anode (half reaction)

TOPICAL QUESTIONS ON EXTRACTION OF METALS

FORM THREE CHEMISTRY.

SECTION A. 20 MARKS.

1.MULTIPLE CHOICE QUESTIONS

In the reaction;

FeCl2 + Cl2 2FeCl3

Chlorine may be regarded as;

A reducing agent

An oxidizing agent

A catalyst

Halogens

ii) The best method to extract iron from its ore is;

Reduction using carbon

Electrolysis of its solution

Roasting of iron ore in air

Electrolysis of its molten form

iii) In the blast furnace, iron ore is converted into iron metal, which of the following statements about the blast furnace is not true;

The reaction in the blast furnace produces a lot of heat which keeps the iron molten

Limestone is reacted with impurities to form slag

Coke is reacted to reduce iron oxide to iron metal

Coke combines with some impurities in the ore

iv) A less electropositive metal may be extracted by;

Ore concentration and electrolytic reduction

Ore concentration and chemical reduction

Ore concentration and purification

Ore concentration, reduction and purification.

v) The reason why carbon is put on the reactivity series of metal is because;

A. Has both metallic and non metallic property

B. They are very electropositive

C. They are reducing agents just like metals

D. They are metallic in nature

vi) which of the following is the most electronegative element?

Sodium

Sulphur

Fluorine

Chlorine

vii) the addition of calcium chloride in sodium chloride during the extraction of sodium helps to;

Increase the solubility of sodium

Lower boiling point of sodium

Increase boiling point of sodium chloride

Make the solution more soluble.

viii) which of the following is not a method of purifying an ore after extraction?

Distillation

Oxidation

Concentration

Electro-refining

ix) the common ore used during the extraction of alluminium?

Bauxite

Cryolite

Feldspar

Kaolin

x) The following are the three steps which are followed during the extraction of any metal except one, which one?

A. Ore concentration

B. Ore purification

C. Refining

D. Extraction

2. Matching items questions.

LIST A

LIST B

An element that reacts by electron loss

Mineral that contain large proportions of metallic or non metallic element

Removing of impurities from the ore without decomposing the ore chemically

Heating ore in absence of air below its melting point

Heating ore in presence of air below its melting point

The apparatus used in extraction of sodium

Compound formed when calcium hydroxide reacts with silica

Iron that contains 4%carbon with other impurities

Method used in extraction of relatively reactive elements.

An apparatus used in the extraction of iron.

Blast furnace

Iron blast

Haematite

Ore

Mineral

Roasting

Desiccating

Metal

Non-metal

Smelting

Out reduction

Electrolysis

Chemical reduction

Dressing

Calcination

Pig iron

Wrought iron

Slag

Limestone

Down cell

Hoffman kiln

(a) (i) Name four important ores of iron.

(ii) What objective is achieved when concentrated ore iron is roasted?

(b) Briefly describe balanced chemical reactions taking place in the blast furnace for the extraction of iron.

(a) List down four (4) common stages in the extraction of less reactive metals like zinc and copper.

Name the ore commonly used in the extraction of iron metal.

The following are series of chemical reactions which occur in the blast furnace during the process of extraction of iron metal.

C(s) + 02→ CO3 + heat (1)

CO, +C→2C0 (2)

Fe203 +3C→2Fe +3C0 (3)

Fe2O3 +CO → Fe +2CO2 (4)

Indicate the two reducing agents in the blast furnace.

Explain the importance of steps (1) to (3).

In this process, a compound "L" which produces a chemical substance that removes impurities as slag is added. Give the name of the substance.

Write the complete chemical reactions that compound "L undergoes to form slag.

study the chart below and answer the questions that follow;

Ore T

StepI Excess sodium hydroxide

solution

residue

Pure aluminium oxide

Cryolite

Na3Al6

Molten aluminium

Molten mixture

Step II

Step III

ElectrolysiStep IV

Name the ore labeled T

State why the ore is first dissolved in excess sodium hydroxide

Write downthe formulaof the aluminium compound in the solution

Namethe process that takes place in step II

Write down the ionic equation taking place in step IV

State why the anode electrode is replaced from time to time

Name one compound present in the residue

Explain why alloys of aluminum are preferred to aluminum for use in overhead power cables.

The diagram below shows the blast furnace used in the extraction of iron. Study it and answer the questions that follows;

ixture of coke limestone and iron Oxide A

Hot air blast

Slag

iron

State the use of the following in the process

Coke

Limestone

Hot air blast

b) (i) state one important use of the pipe marked

(ii) Give the uses of slag in the furnace and as a product of the process

Sodium is extracted by use of a down cell. Study diagram below and answer the questions that follow;

What is a down cell?

Name substance X and Y

In the extraction of sodium, molten calcium chloride is added to the ore, why is this done?

Write equation for the reaction occurring at the;

Anode

Cathode

Sodium is collected at the top most part; explain why this is possible

Mention one environmental effect of extraction of sodium

Explain the role of steel gauze diaphragm

How can we store sodium in laboratory

Mention some of the uses of sodium.

8. (a) what is a blast furnace?

b) Name the common ore used in the extraction of iron

c) One of the reducing agents in blast furnace is limestone, explain two roles of limestone

d) What is steel? Name two uses of steel

e) is there any environmental impact of extraction of iron?

9. ( a) define the term ore?

b) Name the common chief ores of the following metals;

i) Copper

ii) Iron

Aluminium

Zinc

c) Give the factors that determine the methods used in extraction of metals

d) Explain why it took so long for chemists to be able to extract aluminum from its ore

SAMPLE NECTA QN 9 2005

A blast furnace is used to convert iron ore, Fe203 into iron metal. The following is a diagram of the blast furnace for the extraction of iron.

(a) Write down the names of the:

(1) Substances represented by materials V, W, and X.

(ii) Products Y and Z.

(i) Write a balanced chemical equation for the reaction between the iron ore, Fe203 and carbon monoxide.

(ii) What is the function of carbon monoxide in the reaction with Fe203?

If 80kg of iron ore, Fe203 were allowed to react with carbon monoxide during the extraction process, how many kilograms of iron (Fe), would be obtained?

2007

QN. 10. (a) Name the ore commonly used in the extraction of copper metal.

(b) Steps (i) to (iv) below are used during the extraction of copper metal from its ore. Write a balanced chemical equation for each step.

(1) Roasting of the concentrated ore (Cu FeS2) in air.

Heating the roasted ore with silica in the absence of air.

Burning copper sulphide in regulated supply of air.

Purification of copper by electrolysis using copper sulphate solution electrolyte, pure copper cathode and impure copper obtained from the extraction anode.

An iron earring dropped into a container of copper sulphate solution.

Copper knife dipped into zinc nitrate solution.

Copper turnings dropped into a container of dilute hydrochloric acid.

2008.

QN. 9. (a) List down four (4) common stages in the extraction of less reactive metals like zinc and copper.

Name the ore commonly used in the extraction of iron metal.

The following are series of chemical reactions which occur in the blast furnace during the process of extraction of iron metal.

C(s) + 02→ CO3 + heat (1)

CO, +C→2C0 (2)

Fe203 +3C→2Fe +3C0 (3)

Fe2O3 +CO → Fe +2CO2 (4)

Indicate the two reducing agents in the blast furnace.

Explain the importance of steps (1) to (3).

In this process, a compound "L" which produces a chemical substance that removes impurities as slag is added. Give the name of the substance.

Write the complete chemical reactions that compound "L undergoes to form slag.

2011.

QNS. 8. (a) Briefly describe how sodium is extracted in Downs cell. Write all the necessary equations.

(b) List at least four uses of sulphur.

2015.

QN. Describe the extraction of iron from the haematite ore and write all the chemical equations for the reactions involved in each stage of extraction.

2016.

QN8. (a) Identify and state the environmental problem caused by the gas which is released from the blast furnace in the extraction of iron from its oxide

Draw a labelled diagram of a simple electrolytic cell which shows how copper is purified.

Write balanced ionic equations to show the electrode reactions which occur when copper is purified.

2017.

QNS 4. (a) State four steps employed in the extraction of moderate reactive metals. (b) Write balanced equations to show how chlorine reacts with the following:

Water

Aqueous iron (II) chloride solution

Hydrogen sulphide

TOPIC : 9 EXTRACTION OF METALS

School Base-Online

CHEMISTRY EXAMINATION FORM THREE

TOPICAL EXAMINATIONS.

EXTRACTION OF METAL

NAME………………………………………..CLASS…………………………………………….……………TIME: 21/2HRS

INSTRUCTIONS:-

This paper consists of sections A, B and C

Answer all questions

All answers must be written in the spaces provided

All writings should be in blue/black inks except for drawings that should be in pencils

1.MULTIPLE CHOICE QUESTIONS

In the reaction;

FeCl2 + Cl2 2FeCl3

Chlorine may be regarded as;

A reducing agent

An oxidizing agent

A catalyst

Halogens

ii) The best method to extract iron from its ore is;

Reduction using carbon

Electrolysis of its solution

Roasting of iron ore in air

Electrolysis of its molten form

iii) In the blast furnace, iron ore is converted into iron metal, which of the following statements about the blast furnace is not true;

The reaction in the blast furnace produces a lot of heat which keeps the iron molten

Limestone is reacted with impurities to form slag

Coke is reacted to reduce iron oxide to iron metal

Coke combines with some impurities in the ore

iv) A less electropositive metal may be extracted by;

Ore concentration and electrolytic reduction

Ore concentration and chemical reduction

Ore concentration and purification

Ore concentration, reduction and purification.

v) The reason why carbon is put on the reactivity series of metal is because;

A. Has both metallic and non metallic property

B. They are very electropositive

C. They are reducing agents just like metals

D. They are metallic in nature

vi) which of the following is the most electronegative element?

Sodium

Sulphur

Fluorine

Chlorine

vii) the addition of calcium chloride in sodium chloride during the extraction of sodium helps to;

Increase the solubility of sodium

Lower boiling point of sodium

Increase boiling point of sodium chloride

Make the solution more soluble.

viii) which of the following is not a method of purifying an ore after extraction?

Distillation

Oxidation

Concentration

Electro-refining

ix) the common ore used during the extraction of alluminium?

Bauxite

Cryolite

Feldspar

Kaolin

x) The following are the three steps which are followed during the extraction of any metal except one, which one?

A. Ore concentration

B. Ore purification

C. Refining

D. Extraction

2. Matching items questions.

LIST A

LIST B

An element that reacts by electron loss

Mineral that contain large proportions of metallic or non metallic element

Removing of impurities from the ore without decomposing the ore chemically

Heating ore in absence of air below its melting point

Heating ore in presence of air below its melting point

The apparatus used in extraction of sodium

Compound formed when calcium hydroxide reacts with silica

Iron that contains 4%carbon with other impurities

Method used in extraction of relatively reactive elements.

An apparatus used in the extraction of iron.

Blast furnace

Iron blast

Haematite

Ore

Mineral

Roasting

Desiccating

Metal

Non-metal

Smelting

Out reduction

Electrolysis

Chemical reduction

Dressing

Calcination

Pig iron

Wrought iron

Slag

Limestone

Down cell

Hoffman kiln

(a) (i) Name four important ores of iron.

(ii) What objective is achieved when concentrated ore iron is roasted?

(b) Briefly describe balanced chemical reactions taking place in the blast furnace for the extraction of iron.

(a) List down four (4) common stages in the extraction of less reactive metals like zinc and copper.

Name the ore commonly used in the extraction of iron metal.

The following are series of chemical reactions which occur in the blast furnace during the process of extraction of iron metal.

C(s) + 02→ CO3 + heat (1)

CO, +C→2C0 (2)

Fe203 +3C→2Fe +3C0 (3)

Fe2O3 +CO → Fe +2CO2 (4)

Indicate the two reducing agents in the blast furnace.

Explain the importance of steps (1) to (3).

In this process, a compound L which produces a chemical substance that removes impurities as slag is added. Give the name of the substance.

Write the complete chemical reactions that compound L undergoes to form slag.

study the chart below and answer the questions that follows

Ore T

StepI Excess sodium hydroxide

solution

residue

Pure aluminium oxide

Cryolite

Na3Al6

Molten aluminium

Molten mixture

Step II

Step III

ElectrolysiStep IV

Name the ore labeled T

State why the ore is first dissolved in excess sodium hydroxide

Write down the formula of the aluminium compound in the solution

Name the process that takes place in step II

Write down the ionic equation taking place in step IV

State why the anode electrode is replaced from time to time

Name one compound present in the residue

Explain why alloys of aluminum are preferred to aluminum for use in overhead power cables.

The diagram below shows the blast furnace used in the extraction of iron. Study it and answer the questions that follows;

ixture of coke limestone and iron Oxide A

Hot air blast

Slag iron

State the use of the following in the process

Coke

Limestone

Hot air blast

b) (i) state one important use of the pipe marked

(ii) Give the uses of slag in the furnace and as a product of the process

Sodium is extracted by use of a down cell. Study diagram below and answer the questions that follow;

What is a down cell?

Name substance X and Y

In the extraction of sodium, molten calcium chloride is added to the ore, why is this done?

Write equation for the reaction occurring at the;

Anode

Cathode

Sodium is collected at the top most part; explain why this is possible

Mention one environmental effect of extraction of sodium

Explain the role of steel gauze diaphragm

How can we store sodium in laboratory

Mention some of the uses of sodium.

8. (a) what is a blast furnace?

b) Name the common ore used in the extraction of iron

c) One of the reducing agents in blast furnace is limestone, explain two roles of limestone

d) What is steel? name two uses of steel

e) is there any environmental impact of extraction of iron?

9. ( a) define the term ore?

b) Name the common chief ores of the following metals;

i) Copper

ii) Iron

Aluminium

Zinc

c) Give the factors that determine the methods used in extraction of metals

d) Explain why it took so long for chemists to be able to extract aluminum from its ore

School Base-Online Page 6

TOPIC : 9 COMPOUNDS OF METALS

CHAPTER 9

COMPOUNDS OF METALS.

KEY CONCEPTS AND MEANING OF TERMS.

A compound is a substance formed between two or more elements, through chemical means.

An oxide is a compound of an element and oxygen

Direct preparation of oxides involves direct combination of oxygen and a metal.

Soluble oxide can readily dissolve in water example include sodium oxide

Insoluble oxides do not dissolve in water, example is copper oxide

Hydroxide is any inorganic compound that contains hydroxyl group –OH

Amphoteric hydroxides can react with both acids and bases.

Metal carbonate- this is a compound of a metal combined with a carbonate radical

Test for metal carbonate- All metal carbonates reacts with dilute nitric acids to form a salt, water and carbon dioxide.

Nitrates- these are salts derived from nitric acid.

Brown ring test- this is a technique used to test for presence of a nitrate in a substance

All nitrates are soluble in water

The behavior of a nitrate when heated depends on the position of the metal in reactivity series

Chloride do not decompose on heating a part from ammonium chloride

When concentrated sulphuric acid is added to a compound and hydrogen chloride gas is produced, this confirms the compound is a chloride.

Hydrochloride gas reacts with ammonia gas to form white fumes of ammonia chloride.

Hydrated zinc sulphate is light green in colour; when heated, it loses water of crystallization to become white.

Barium and lead ions are used to test for the presence of sulphate in a compound.

HOW TO TACKLE QUESTIONS IN TOPIC.

A Student who passes this paper should be able to;

Write correct formula of oxides,hydroxides, carbonates, hydrogencarbonates, sulphates, chlorides, and nitrates of various metals.

Outline a simple procedure on how to prepare samples of above compounds.

State how each of the above compounds can be tested in laboratory

With equations, explain what happens on the above compounds when heated.

Account for solubility of the above compounds in water.

Identify the various products formed when above compounds are heated

Write equations for the formations of various compounds

Identify the unique colors displayed by various compounds both when hydrated and when dry.

INTRODUCTION.

A compound is a substance formed by the chemical combination of two or more elements. Metals react with nonmetals to form various compounds. The common compounds formed by metals include: Oxides, hydroxides, carbonates, sulphates, nitrates, chlorides and hydrogen carbonates.

Metal Oxides

A metal oxide is a chemical compound which is formed when a metal atom combines with oxygen. The anion of a metal oxide is in the form of oxide ion (02). Examples of metal oxides include:

Potassium oxide (K20),

Sodium oxide (NaO),

Calcium oxide (CaO),

Magnesium oxide (MgO),

Aluminium oxide (Al203),

Zinc oxide (ZnO),

Iron (III) oxide (Fe203).

Metal oxides are binary compounds, which means they are made up of only two elements.

Preparation of oxides

Metal oxides can be prepared by either the direct method or by the indirect method.

Preparation of metal oxide by direct method.

The direct method involves direct heating of a metal in oxygen to yield an oxide while the indirect method involves the decomposition of a metal compound by heating to give an oxide. Preparation of metal oxide by direct methods can be done indifferent ways.

1. By direct heating of element with oxygen

In this method, a metal is directly heated in oxygen to give a metal oxide. Many metals burn rapidly when heated in oxygen or in air, hence producing their oxides.

The following oxides can be prepared by this method

Sodium Oxide.

Sodium burns in air with a persistent yellow flame to form sodium oxide which is a white solid.

In excess air, sodium burns to form sodium peroxide which is a pale yellow solid.

Sodium + Oxygen

Sodium peroxide

2Na(s) + 02(g)

Na202(s)

When sodium dissolves in water at room temperature, it forms sodium hydroxide which is an alkaline solution.

When sodium peroxide is dissolved in ice cold water, hydrogen peroxide is formed.

Magnesium

Magnesium burns in oxygen with a brilliant white flame to form a white solid of magnesium oxide. The oxide reacts with water to form basic magnesium hydroxide

Magnesium + Oxygen

Magnesium oxide

2Mg(s) + 02(g )

2MgO

Figure 7.1 Magnesium burning in oxygen.

Calcium

Calcium burns in air with a red flame to form calcium oxide which is a white solid.

Calcium + Oxygen

Calcium oxide

2Ca(s) + 02(g)

2CaO(s)

When calcium hydroxide is dissolved in water, basic hydroxide is formed. Calcium hydroxide is however sparingly soluble in water.

The reaction of metals with oxygen depends on the position of the metal in reactivity series. Potassium and sodium react vigorously with oxygen to form metal oxides.

Preparation of oxides by indirect method

This involves thermal decomposition of certain compounds like hydroxides, carbonates and nitrates. The carbonates of calcium, zinc, lead, copper and iron (III) decompose on heating to give a metal oxide and carbon (IV) oxide gas. First, the metal compound is prepared and then decomposed to form metal oxides

Calcium carbonate

Calcium oxide + Carbon (IV) oxide

CaCO3(s)

CaO(s) + CO2 (g)

Copper (II) carbonate

Copper (II) oxide + Carbon (IV) oxide

CuCO3(s)

CuO(s) + CO2(g)

Zinc carbonate

Zinc oxide + Carbon(IV) oxide

ZnCO3(s)

ZnO(s) + CO2(s)

Lead(II) carbonate

Lead(II) oxide + Carbon(IV) oxide

PbCO3(s)

PbO(s) + CO2(s)

Nitrates of average reactivity decompose on heating to yield an oxide, nitrogen dioxide gas and oxygen gas.

Calcium nitrate

Calcium oxide +Nitrogen(IV) oxide + Oxygen

2Ca(NO3)2(s)

2CaO(s) + 4NO2(g) +02(g)

(white) (white) brown gas colourless

2Zn(NO3)2(s)

2ZnO(s) + 4NO2(g) +02(g)

(white) (yellow when hot and white whencold)

Lead nitrate

Lead(II) oxide +Nitrogen(IV) oxide + Oxygen

2Pb(NO3)2(s)

. 2PbO(s) + 4NO2(g) + 02(g)

(white) (Orange when hot and yellow when cold)

Copper(II) nitrate

Copper(II) oxide +Nitrogen(IV) oxide + Oxygen

2 Cu(NO3)2

2CuO(s) + 4NO2(g) + 02(g)

(Black)

3. Preparation of oxides by oxidation of some metals with nitric(V) acid

Concentrated nitric (IV) acid is a powerful oxidizing agent. It readily oxidizes metals while itself its reduced into nitrogen (II) oxide, water and oxygen.

Copper + Nitric (V) acid →Copper (II)oxide + Nitrogen(IV) oxide + Water

2Cu(s) + 8HNO3→2CuO (aq) + 8N0(g) +4H20(1) + 02(g)

Magnesium burns in oxygen with a white flame to form a white solid while calcium burns in oxygen with a red flame and forms a white solid.

When calcium carbonate is heated, a white solid is formed and a colourless gas is evolved.

When zinc hydroxide is heated a white solid is formed and colourless liquid droplets seen.

Carbonates decompose on heating to form solid oxides and carbon (IV) oxide.

Hydroxides decompose on heating to form respective oxides and water.

Very reactive metals such as sodium, potassium, magnesium and calcium readily react with oxygen even without heating to form their oxides. This explains why sodium and potassium are stored under paraffin. Magnesium and calcium form oxide coatings around them, thus before reacting them with any substance, they have to be scrubbed to remove the coating.

At high temperatures, oxygen reacts with many compounds forming oxides, e.g

Sulphides are usually oxidized when heated with oxygen.

Lead sulphide + Oxygen→Lead oxide + Sulphur(IV) oxide

PbS(s) + 302(g) →2PbO(s) + 2502(g)

Zinc sulphide + Oxygen → Zinc oxide + Sulphur(IV) oxide

2ZnS(s) + 302(g) → 2ZnO(s) + 2S02(g)

A black solid is observed at the bottom of the test tube and red brown fumes are observed at the mouth of the test tube.

Concentrated nitric(V) acid oxidisesmetals while itself its reduced to nitrogen(IV) oxide, water and oxygen. Copper is oxidised to copper(III)oxide.

Classification of metal oxides

Metal oxides are classified into various types depending on their solubility in water and their reactions with acids.

Classification based on solubility

Soluble metal oxidesreadily dissolve in water to form solutions. A soluble metal oxide is known as alkali. Examples of soluble metal oxides include: potassium oxide, sodium oxide and calcium oxide.

Amphoteric oxides: These are metal oxide which reacts with both acid and bases. Lead oxide, copper(II) oxide, aluminium oxide, zinc oxide and iron (III) oxide are all amphoteric oxides. When Amphoteric oxide react with alkalis, complex bases are formed.

Properties of metal oxides

Reaction with water

Metal oxides react differently with water. Most metal oxides are insoluble in water.

Sodium and potassium oxide are highly soluble in water.

Sodium oxide reacts with water to give sodium hydroxide which is a strong base.

Sodium oxide + Water → Sodium hydroxide

Na20(s) + H20(1) → 2NaOH(aq)

Magnesium oxide reacts with water to form basic magnesium hydroxide solution.

Magnesium oxide + Water →Magnesium hydroxide

MgO(s) + H20(1) → Mg(OH)2(aq)

Calcium oxide is sparingly soluble in water. It dissolves in water to form calcium hydroxide.

Calcium oxide + Water → Calcium hydroxide

The solutions formed are alkaline in nature and turns litmus paper blue

Reactions of oxides with acids.

Most metal oxides reacts with acids to form salts and water, this reaction is called neutralization.

Dilute hydrochloric acid reacts with metal oxides to form chlorides and water.

Sodium oxide + Hydrochloric acid →Sodium chloride + Water

Na20(s) + 2HC1(aq) → 2NaC1(aq) + H20(1)

Calcium oxide + Hydrochloric acid →Calcium chloride + Water

CaO(s) + 2HC1(aq) → CaC12(aq) + H20(1)

The reaction between lead (II) oxide and dilute hydrochloric acid stops as soon as it is formed. This is due to the formation of lead (II) chloride which is insoluble.

Lead(II) oxide + Hydrochloric acid →Lead(II) chloride + Water

PbO(s) + 2HC1(aq) → PbC12(s) + H20(1)

The lead (II) chloride forms a coating around the lead (II) oxide preventing further action of the acid on the oxide.

Dilute nitric(V) acid reacts with metal oxides to yield a nitrate and water.

Sodium oxide + Nitric(V) acid →Sodium nitrate + Water

Na20(s) + HNO3(aq) → NaNO3(aq) +H20(1)

Calcium oxide + Nitric(V) acid → Calcium nitrate + Water

Magnesium,Zinc, Aluminium and lead are basic oxides. They react with nitric acid to form salt and water only. Aluminium,zincandleadoxide reacts with sodium hydroxide and nitric acids and thus they are amphoteric.

Dilute sulphuric acid and hydrochloric acid cannot be used with lead II Oxide salts because, they form insoluble oxides.

The insoluble oxides forms a coating on the oxides preventing any further reaction.

When lead(II) oxide is added in water, it does not dissolve while other oxides namely calcium oxide and sodium oxides dissolves in water to form a colorlesssolutions.

Uses of metal oxides

Calcium oxide is used in preparation of calcium carbide which is used to make fertilizer, in steel making and in carbide lamp.

CaO(s) + 3C(s)→CaC2(s) + CO2(g)

Calcium oxide is used in lining furnaces since it can withstand high temperatures.

In the extraction of iron, calcium carbonate is added into the blast furnace to remove impurities from the iron and form slag.

Calcium oxide is used as a drying agent in the preparation of ammonia gas and ethanol.

Calcium oxide is used in the preparation of mortar. Calcium oxide is reacted with water to make slaked lime (Ca(OH)2). Slaked lime, sand and water are mixed to make a solution which is used to stick bricks together.

In the manufacture of cement, clay (aluminium silicate) is heated with calcium oxide (lime) to form cement. Cement is a mixture of calcium silicates and aluminates.

Soil treatment: Calcium oxide (lime) is added to soils with a low pH(acidic soils) to reduce the acidity.

Magnesium oxide is used in lining of steel furnaces because of its high melting point.

Zinc oxide is used as a white pigment in paints, a filter in rubber and as a component in enamel and antiseptic ailments.

Naturally occurringaluminium oxide is used as an abrasive.

Hydroxides

A metal hydroxide is any inorganiccompound that contains the hydroxyl group –OH

Most metal hydroxides are bases and have a PH greater than 7. They neutralize acids. Hydroxides of alkali metals such as sodium and potassium are soluble in water forming very strong alkalis, while those of magnesium and calcium are sparingly soluble in water. Like oxides, hydroxides can be prepared by direct and indirect methods. Direct method is suitable for the most reactive elements.

i) Preparation of metal hydroxides by direct method

Metals high in reactivity (such as potassium, sodium, calcium and magnesium) react with water to form the hydroxide.

Preparation of sodium hydroxide

When a piece of sodium metal is dropped into cold water, it darts about with a hissing sound as the metal piece is slowly consumed. Eventually, the sodium piece disappears as it has reacted with the water to form sodium hydroxide while liberating hydrogen gas.

Sodium +Water →Sodium hydroxide + Hydrogen

2Na(s) + 2H20(1) → 2NaOH(aq) + H2(g)

The other highly reactive metals react with cold water producing their corresponding hydroxides and liberate hydrogen gas.

Potassium +Water →Potassium hydroxide + Hydrogen gas

K(s) + H20(1) → KOH(aq) + H2(g)

Potassium reacts vigorously with water.

Calcium +Water → Calcium hydroxide + Hydrogen gas

Ca(s) + 2H20(1) → Ca(OH)2(aq) + 2H2(g)

Calcium hydroxide is alkaline and sparingly soluble in water.

Magnesium +Water →Magnesiumhydroxide + Hydrogen gas

Mg(s) + 2H20(1) → Mg(OH)2(aq) + H2(g)

This reaction is much slow than that of calcium.

Preparation of metal hydroxide by indirect method

This involves reactions between a salt solution and alkalis. The metal salt chosen determines the metal hydroxide to be formed. The alkalis displace the metals from their salts enabling those metals to combine with the hydroxyl ion.

Examples

Zinc chloride + Sodium hydroxide→Zinc hydroxide + Sodium chloride

ZnC12(aq) + 2NaOH(aq)→Zn(OH)2 (aq) + 2NaCl(aq)

Copper(II) sulphate + Sodium hydroxide→Copper(II) hydroxide + Sodium sulphate

CuSO4(aq) + 2NaOH(aq)→Cu(OH)2(aq) + Na2SO4(aq)

Iron(III) chloride + Sodium hydroxide→Iron(III) hydroxide + Sodium chloride

FeC13(aq) + 3NaOH(aq)→Fe(OH)3(aq) + 3NaC1(aq)

The method is suitable in the preparation of hydroxides of the metals with average reactivity.

When calcium is added on water, a colourless gas is evolved and a solution which turns red litmus paper to blue is formed.

Calcium reacts with water to form calcium hydroxide solution and hydrogen gas is evolved.

When iron(III) chloride is added to sodium hydroxide a red-brown solid is observed.

An insoluble hydroxide is formed by the reaction of iron(III) chloride and sodium hydroxide.

Iron(III) chloride + Sodium hydroxide→Sodium chloride + Iron (III) hydroxide

FeC13(aq) + NaOH(aq) →NaCl(aq) + Fe(OH)3(aq)(red-brown)

Classification of hydroxides

Just like metal oxides, metal hydroxides are clarified according to their solubility in water and according to their reaction with acids.

Classification based on solubility

soluble hydroxides

insoluble hydroxides

Soluble hydroxides are soluble in water. Hydroxides of potassium and sodium are highly soluble in water. Magnesium and calcium hydroxides are sparingly soluble in water. All the other metal hydroxides are insoluble in water.

Classification based on reaction with acids

Basic hydroxides

Amphoteric hydroxide

All metal hydroxides are basic in nature. They all react with acids to form salt and water only.

Zinc hydroxide + Hydrochloric acid →Zinc chloride + Water

Zn(OH)2(s) + 2HC1(aq)→ZnC12(aq) + 2H20(1)

Aluminium hydroxide + Sulphuric(VI) acid→Aluminiumsulphate + Water

2A1(OH)3(s) + 3H2SO4(aq)→Al2(SO4)3(aq) + 6H20(1)

Lead(II) hydroxide + Nitric(V) acid→Lead(II) nitrate + Water

Pb(OH)2(s) + 2HNO3(aq) →Pb(NO3)2(aq) + 2H20(1)

Iron(III) hydroxide + Sulphuric(VI) acid→Iron(III) sulphate + Water

2Fe(OH)3(s) + 3H2SO4(aq)→Fe2(SO4)3(aq) + 6H20(1)

Amphoteric metal hydroxides react with both bases and acids. They are insoluble hydroxides. The amphoteric hydroxides are: lead hydroxide (Pb(OH)2, aluminiumhydroxide(Al(OH)3) and zinc hydroxide (Zn(OH)2).

The reactions with acids are known as neutralization reactions.

When they react with bases, they form complex ions . When they react with excess sodium hydroxide, the insoluble amphoteric hydroxides dissolve to form complex ions.

Their reactions are as shown,

Lead hydroxide + sodium hydroxide →sodium plumbate + water

Pb(OH)2(s) + 2NaOH(aq)→Na2PbO2(aq) + 2H20(1)

Zinc hydroxide + sodium hydroxide→sodium zincate + water

Zn(OH)2(s) + 2NaOH(aq) →Na2Zn02(aq) + 2H20(1)

Aluminium hydroxide + sodium hydroxide →sodium aluminate + water

2A1(OH)3(s) + 2NaOH(aq) →2NaA103(aq) + 2H20(1)

Chemical properties of metal hydroxides

Action on heating

Hydroxides of potassium, sodium and calcium do not decompose on heating.

Hydroxides of metals lower in reactivity decompose on heating to form metal oxide and water.

Lead(II) hydroxide

Lead(II) oxide + Water

Pb(OH)2

Pb0(s) + H20(1)

Zinc hydroxide

Zinc oxide + Water

Zn(OH)2

ZnO(s) + H20(1)

Iron(III) hydroxide

Iron(III) oxide + Water

Action on acid

Metal hydroxides react with mineral acids to give salt and water.

Sodium hydroxide + Hydrochloric acid

Sodium chloride + Water

NaOH(aq) + HCI(aq)

NaCl(aq) + H20(1)

Zinc hydroxide + Hydrochloric acid

Zinc chloride + Water

Zn(OH),(s) + 2HC1(aq)

ZnC12(aq) + 2H20(1)

Aluminium hydroxide + Sulphuric(VI) acid Aluminiumsulphate + Water

2A1(OH)3(s) + H2SO4(aq)

Al2(SO4)3(aq) + 2H20(1)

Lead(II) hydroxide + Nitric(V) acid

Lead(II) nitrate + Water

Pb(OH)2(s) + 2HNO3(aq)

Pb (NO3)2( aq) + 2H20 (1)

Iron(III) hydroxide + Sulphuric(VI) acid

Iron(III) sulphate + Water

2Fe(OH)3(s) + H2SO4(aq)

Fe2(SO4)3(aq) + H20(1)

Uses of metal hydroxides.

Calcium hydroxide is used in soil treatment

Aluminium and magnesium hydroxides are used as antacids to neutralize stomach acids.

Slaked lime is used in building, to smear walls and also in preparation of mortar

Calcium hydrogensulphate is used in preparing sulphite pulp from wood, for making paper and artificial silk

Calciumhydroxide is used to remove temporary hardness of water

Calcium hydroxide is used in the manufacture of undercoat paints, which are applied first before applying gloss paints.

Calcium hydroxide is used in qualitative analysis, for example testing for carbon dioxide.

Sodium hydroxide is used in the extraction of metal

Carbonates and hydrogen carbonates

They are derived from carbonic acid.

The hydrogen ions in carbonic acid (H2CO3) are replaced by a metal. Carbonic acid is dibasic which means it has two replaceable hydrogen ions per molecule. If both hydrogen ions are replaced, a carbonate is formed, but if only one of the hydrogen ions is replaced then a hydrogen carbonate is formed.

Preparation of metal carbonates

They are formed when both hydrogen ions are replaced by a metal ion.

The methods used to prepare carbonates depend on whether the carbonate to be prepared is soluble or insoluble in water.

i) Preparation of soluble metal carbonates

Sodium, potassium and ammonium carbonates are the only soluble metal carbonates.

They are prepared by reacting the corresponding alkali with carbon(IV) oxide.

Alkali + Carbon(IV) oxide→Metal carbonate + water.

Sodium carbonate is also called washing soda.

Sodium hydroxide + Carbon(IV) oxide →Sodium carbonate + Water

2NaOH(aq) + CO2(g) →Na2CO3(aq) + H20(l)

Potassium hydroxide + Carbon(IV) oxide →Potassium carbonate + Water

ii) Preparation of insoluble metal carbonates

All insoluble metal carbonates are prepared by precipitation method. A soluble carbonate is reacted with a soluble salt which contains the desired cations.

Example

To prepare calcium carbonate, a soluble carbonate such as sodium carbonate is reacted with any soluble salt of calcium such as calcium chloride. The salts interchange their anions thus producing the desired salt as a precipitate (an insoluble solid) and another salt as a solution. The mixture is filtered then the desired salt calcium carbonate is obtained as a residue. It is then rinsed using distilled water to remove impurities and dried.

Calcium chloride + Sodium carbonate→Calcium carbonate + Sodium chloride

CaC12(aq) + Na2CO3(aq)→CaCO3(s) + 2NaC1(aq)

Ionic equation:

Ca2+(aq) + 2C1-(aq) + 2Na+(aq) + CO2-(aq) →CaCO3(s) + 2Na+(aq) + 2C1-(aq)

Net ionic equation:

Ca2+(aq) + CO2-(aq)→ CaCO3(s)

Similarly;

When soluble carbonates react with magnesium sulphate solution which is soluble in water, a white precipitate of insoluble magnesium carbonate is formed.

Sodium carbonate + Magnesium sulphate→Sodiumsulphate + Magnesium carbonate

Na2CO3(aq) + MgSO4(aq) →Na2SO4(aq) + MgCO3(s)

When carbon(IV) oxide is passed to a solution of sodium hydroxide, a white solid is observed. It dissolves in the solution.

While on passing carbon(IV) oxide to lime water, a white precipitate is observed.

The precipitate dissolves on excess passage of carbon(IV) oxide to form a colourlesssolution.

On adding sodium carbonate solution to copper(II) sulphate, a bright blue solid is observed.

Carbon(IV) oxide reacts with sodium hydroxide to form sodium carbonate and water.

Sodium hydroxide + Carbon(IV) oxide →Sodium carbonate + Water

NaOH(aq) + CO2(g)→Na2CO3(aq) + H20(1) (White)

Carbon(IV) oxide reacts with calcium hydroxide to form calcium carbonate solid.

The solid dissolves on excess passage of carbon(IV) oxide.

Calcium hydroxide + Carbon(IV) oxide →Calcium carbonate + Water

Ca(OH)2(aq) + CO2(g)→CaCO3(s) + H20(1)

(White)

Calcium carbonate + Water + Carbon(IV) oxide →Calcium hydrogen carbonate

CaCO3(s) + H20(1) + CO2(g)→Ca(HCO3)2(aq) (Colourless solution)

Copper(II) sulphate reacts with sodium carbonate to form copper(II) carbonate and sodium sulphate.

Copper(II) sulphate + Sodium carbonate →Copper(II) carbonate + Sodium sulphate

CuSO4(aq) + Na2CO3(aq)→CuCO3(s) + Na2SO4(aq)

Classification of metal carbonates

Metal carbonates are classified according to their solubility in water. They are classified into soluble carbonates and insoluble carbonates.

Soluble carbonates

These are metal carbonates of sodium, potassium and ammonium. They are soluble in water.

Insoluble carbonates

All carbonates of metals lower than sodium in the reactivity series are insoluble.

Chemical properties of carbonates

Effect of heat

Carbonates of very reactive metals such as sodium and potassium carbonates are stable. They do not decompose on heating.

All carbonates of metals lower than sodium in the reactivity series decompose on heating to give a metal oxide and carbon(IV) oxide gas. Carbon(IV) oxide gas forms a white solid with lime water.

Calcium carbonate →(heat) Calcium oxide + Carbon(IV)oxide

CaCO3→(heat) Ca0(s) + CO,

Zinc Carbonate →(heat)Zinc Oxide + Carbon(IV)oxide

Zn CO3(s) →(heat) ZnO + CO2(g)

Lead carbonate →(heat) lead oxide + Carbon(IV)oxide

PbCO3(s) → PbO(s) + CO2(g)

Copper carbonate →(heat)Copper(II) oxide + Carbon(IV)oxide

CuCO3→(heat)CuO + CO2

Copper(II) carbonate turns to a black solid and a colorless gas is evolved.

Copper(II) carbonate decomposes on heating to form copper(II) oxide, a black solid and carbon(IV) oxide gas.

Copper(II) carbonate

Copper(II) oxide + Carbon(IV) oxide

Action on dilute acids

Metal carbonates react with mineral acids (hydrochloric acid, sulphuric(VI) acid) to give a salt, carbon(IV) oxide and water.

Zinc carbonate + Sulphuric(VI) acid →Zinc sulphate + Water + Carbon(IV) oxide

ZnCO3(s) + H2SO4(aq)→ZnSO4(aq) + H20(1) + CO2(s)

Sodium and zinc carbonates reactwith dilute hydrochloric acid to give colourless chloride solutions carbon dioxide and water.

Na2CO3(s)+ 2HCl (aq) → 2NaCl(aq) + CO2(g) + H2O(l)

ZnCO3(s) + 2HCl (aq) →ZnCl2(aq) +CO2(g) +H2O(l)

The reaction of theacid with copper II carbonate produces carbon dioxide and a blue solution of copper II Chloride

CuCO3(s) + HCl(aq) → PbCl2(s) + CO2(g) + H2O(l).

The lead chloride forms a coating on the lead carbonate preventing further reaction.

Lead carbonate and calcium carbonate reacts with dilute sulphuric acid to form insoluble lead and calcium sulphate respectively.

PbCO3(s) + H2SO4(aq) →PbSO4(s) +CO2(g) + H2O(l)

CaCO(s) + H2SO4(aq) →CaSO4(s) + CO2(s) + H2O(l)

The insoluble sulphates form a coating on the carbonates thus preventing further reactions between carbonates and the acid.

Dilute nitric acid reacts with all the metal carbonates including lead and calcium carbonates to liberate carbondioxide and a soluble metal nitrate.

PbC03(s) + 2HNO3(aq) →Pb(NO3)2(aq) + CO3(g) +H2O(l)

CaCO3(s) + 2HNO3(aq) →Ca(NO3)2(aq)+CO2(g) +H2O(l)

Test for carbonates

Carbonates decompose on heating to produce carbon(IV) oxide which forms a white solid with lime water.

Ca(OH)2(aq) + CO2(g) → CaCO3(s) +H20(1)

lime water white ppt

In excess carbon(IV) oxide white precipitate of calcium carbonate dissolves to form calcium hydrogen carbonate solution.

Calcium carbonate + Water + Carbon(IV) oxide → Calcium hydrogen carbonate

CaCO3(s)+ H20(1) + CO2(g) →Ca(HCO3)2(aq)(Colourlesssolution)

This is the test for carbonates.

Uses of carbonates

Sodium carbonate is used to soften water.

Sodium carbonate is also used in the manufacture of glass.

When sodium carbonate is heated together with sand (silicon(IV) oxide), a concentrated solution of sodium silicate is formed. This sodium silicate is glass.

Sodium carbonate + Silver(IV) oxide →Sodium silicate + Carbon(IV) oxide

Na,CO3 + 2 SiO,→2NaSiO3 +2NaHCO3(aq)

4. Sodium hydrogen carbonate is used in baking and removing grease.

Metal hydrogen carbonates

These are compounds formed when one of the hydrogen ions of carbonic acid is replaced by a metal. All metals form hydrogen carbonates except Aluminium, lead, zinc, iron and copper.

Preparation of metal hydrogen carbonates

They are formed when concentrated hydroxides react with excess carbon (IV) oxide gas.

Preparation of sodium hydrogen carbonate

Its prepared by passing carbon(IV) oxide (CO2) gas through concentrated sodium hydroxide.

The reaction occurs in two stages. First the carbonate is formed which then reacts with excess carbon(IV) oxide to form the hydrogen carbonate.

Stage 1 Formation of sodium carbonate

Sodium hydroxide + Carbon(IV) oxide→Sodium carbonate + Water

2NaOH(aq) + CO2(g) →Na2CO3(aq) + H20(1)

Stage 2 Formation of sodium hydrogen carbonate

The sodium carbonate formed then reacts with excess carbon(IV) oxide to form it.

Sodium carbonate + Carbon(IV) oxide + Water →Sodium hydrogen carbonate

Na2CO3(aq) + CO2(g) + H20(1) → NaHCO3

Chemical properties of metal hydrogen carbonates

Effect on heat

Hydrogen carbonates, when heated, decompose to yield a metal carbonate, carbon(IV) oxide gas and water.

Examples:

Sodium hydrogen carbonate→ Sodiumcarbonate + Carbon(IV) oxide + Water

2NaHCO3→Na2CO3 + CO2 + H2O

Calcium hydrogen carbonate → Calciumcarbonate + Carbon(IV) oxide + Water

Ca(HCO3)2 →CaCO3 + CO2 + H2O

Reaction with acids

Hydrogen carbonates react with acids to yield a salt, carbon(IV) oxide gas and water.

Sodium hydrogen carbonate + Hydrochloric acid→ Sodium chloride +Carbon(IV) oxide + Water

NaHCO3 + HCl→ NaC1 + CO2 + H2O

Test for soluble hydrogen carbonates.

Sodium hydrogen carbonates reacts with magnesium sulphate to form an insoluble salt of magnesium carbonate and sodium sulphate

MgSO4(aq) + Na2CO3(aq) →MgCO3(s) + Na2SO4(aq).

Sodium hydrogen carbonate reacts with magnesium sulphate to form soluble magnesium hydrogen carbonate and sodium sulphate

MgSO4(aq) + 2NaHCO(aq) →Na2SO4(aq) + Mg(HCO3)2 (aq).

Soluble carbonates form a white precipitate with magnesium sulphate solution while hydrogen carbonates do not. This test is used to distinguish soluble carbonates from hydrogen carbonates.

Uses of hydrogen carbonates

I. Sodium hydrogen carbonate is used in preparation of baking powder.

2. Sodium hydrogen carbonate acts as an antacid which neutralizes the hydrochloric acid in the stomach.

3. They can be used in manufacture of glass

Nitrates

Nitrate are salts formed from nitric(V) acid.

Preparation of metal nitrates

There are several methods for preparing metal nitrates. The methods include:

1. Reactions between metals and dilute nitric(V) acid

Nitric(V) acid oxidizes metal to form a metal nitrate and hydrogen gas.

Sodium + Nitric(V) acid→Sodium nitrate + Hydrogen gas

Na(s) + HNO3(aq)→ NaNO3(aq) + H2(g)

Calcium + Nitric(V) acid→Calcium nitrate + Hydrogen gas

Ca(s) + 2 HNO3(aq)→ Ca(NO3)2 + H2(g)

Reaction between metal hydroxide and dilute nitric(V) acid

The reaction involves interchange of ions. It is a neutralization reaction.

Sodium hydroxide + Nitric(V) acid →Sodium nitrate + Water

NaOH(aq) + HNO3(aq)→NaNO3(aq) + H20(1)

Magnesium hydroxide + Nitric(V) acid→Magnesium nitrate + Water

Mg(OH)2(aq) + 2HNO3(aq)→Mg(NO3)2(aq) + H20(1)

Aluminium hydroxide + Nitric(V) acid→Aluminium nitrate + Water

Al(OH)3(aq) + 3HNO3(aq)→Al(NO3)3(aq) + 3 H20(1)

Reaction between metal oxide and dilute nitric(V) acid

This is another neutralization reaction. It leads to formation of a nitrate salt and water.

Copper(II) oxide + Nitric(V) acid→Copper(II) nitrate + Water

CuO(s) + 2HNO3(aq)→Cu(NO3)2(aq) + H20(1)

Calcium oxide + Nitric(V) acid→Calcium nitrate + Water

Reactions between metal carbonates and dilute nitric (V) acid

When carbonates react with nitric(V) acid, the reaction produces a nitrate salt, carbon(IV) oxide gas and water.

Copper carbonate + Nitric(V) acid→Copper nitrate + Carbon(IV) oxide

CuCO3(s) + 2HNO3(aq)→Cu(NO3)2(aq) + CO2(g) + H20(1)

Sodium carbonate + Nitric(V) acid→Sodium nitrate + Carbon(IV) oxide

Na2CO3(s) + 2 HNO3(aq)→2NaNO3(aq) + CO2(g) + H20(1)

Zinc carbonate + Nitric(V) acid →Zinc nitrate + Carbon(IV) oxide

ZnCO3(s) + 2HNO3(aq)→Zn(N0)2(aq) + CO2(g) + H20(1)

5. Reaction between metal hydrogen carbonate and dilute nitric(V) acid

Hydrogen carbonates react with nitric(V) acid in a similar way carbonates does. The products of the reaction are a metal nitrate, carbon dioxide gas and water.

Calcium hydrogen carbonate +Nitric(IV) acid→ Calcium nitrate +Carbon(IV) Oxide + Water

Ca(HCO3)2(1) + 2HNO3(aq)→Ca(N0)2(aq) + 2CO2(g) + 2H20(1)

Sodium hydrogen carbonate +Nitric(IV) acid→ Sodium nitrate +Carbon(IV) Oxide + Water

NaHCO3(1) + HNO3(aq) →NaNO3(aq) + CO2(g) + H20(1)

Chemical properties of metal nitrates

Effect of heat

All metal nitrates decompose on heating. The products of the decomposition depend on the reactivity of the metal forming the nitrate.

2NaNO3(s) → 2NaNO2(s) + 02(g)

Potassium nitrate →potassium nitrite + Oxygen

heat

2KNO3(s) →2KNO2(s) + 02(g)

All nitrates of calcium, aluminium, magnesium, zinc, iron, lead and copper decompose on heating to yield a metal oxide, nitrogen dioxide gas and oxygen gas.

Calcium nitrate →Calcium oxide +Nitrogen(IV) oxide + Oxygen

2Ca(NO3)2(s) →2CaO(s) + 4NO2(g)+ 02(g)

(white) (white) red-brown gas colourless

Zinc nitrate → Zinc oxide + Nitrogen(IV)oxide + Oxygen

2Zn(NO3)2(s) →2ZnO(s) + 4NO2(g)+ 02(g)

(white) (yellow when hot and white when cold)

Lead(II) nitrate →Lead(II) oxide +Nitrogen(IV) oxide + Oxygen

2Pb(NO3)2(s) →2Pb0(s) + 4NO2(g)+ 02(g)

(white) (Orange when hot and yellowwhen cold)

Copper(II) nitrate → Copper(II) oxide + Nitrogen(IV) oxide + Oxygen

2Cu(NO3)2→2CuO(s) + 4NO2(g) + 02(g) (blue) (black)

Silver nitrate →Silver + Nitrogen(IV) oxide + Oxygen

The nitrates of silver (Ag), mercury (Hg) and gold (Au) decompose to form metal, brown gas (NO2) and oxygen gas.

2AgNO3→2Ag(s) + 2NO2(g) + 02(g)

Brown gas.

Mercury nitrate →Mercury + Nitrogen(IV) oxide + Oxygen

Hg(NO3)2→Hg(s) + 2NO2(g) + 02(g)

Gold nitrate →Gold + Nitrogen(IV) oxide + Oxygen

2AuNO3(s) → 2Au(s) + 2NO2(g) + 02(g)

Test for nitrates

Various tests can be used to identify metal nitrates in solution or in solid form .these tests include action of heat on nitrates. When any metal nitrates is heated, reddish brown fumes of nitrogen dioxide are observed. Other tests includes the brown ring test and the use of copper.

The brown ring test.

When concentrated sulphuric acid is poured in a test tube with test solution and iron II sulphate solution, it moves and settles at the bottom because it is denser. A brown ring forms between the layer of concentrated sulphuric acid and the mixtureof the solution and iron II sulphate, if the test solution is a nitrate.

All nitrates form this brown ring, this test for nitrates is therefore called a brown ring test.

The formation of brown ring.

Concentrated sulphuric acid reacts with the nitrate ions to form a nitric acid.

H2SO4(I) + 2NO3-(aq) →2HNO3(aq) + SO42-(aq).

The nitric acid formed oxidizes iron IIsulphate to iron (III).

Fe2+(aq) → Fe(aq) + e-

Nitric acid acid is itself reduced to nitrogen monoxide;

4HNO3(aq) → 4NO(g) + 2H2O(l) +3O2(g)

Nitrogen monoxide combines with some of the remaining iron (II) SULPHATE to form dark brown complex [FeSO.NO(aq)]

FeSO4(aq) + NO(g) → FeSO4.NO(aq).

If the solution is disturbed, the brown ring disappears. This is because concentrated sulphuric acid and water mix, producing a lot of heat, which help to decompose the FeSO4.NO.The FeSO.NO is very unstable compound

FeSO4.NO(aq) → FeSO4(aq) + NO(g)

The formation of the brown complex is a reversible chemical reaction.

FeSO4(aq) + NO(g) FeSO4.NO(aq)

A brown ring is observed when iron(II) sulphate is added to the acidified sample solution.

The ring is also observed when the solid sample is warmed with concentrated sulphuric(VI) acid.

When potassium nitrate is heated it turns to a slight yellow solid and a colorless gas is produced.

Lead (II) nitrate turns to an orange solid, which changes to yellow solid when it has cooled.

Uses of nitrates

In agriculture - Nitrogenous fertilizers are mainly nitrates. They are ammonium nitrate, potassium nitrate and calcium nitrate.

In photography -Silver nitrate solution reacts with potassium bromide to form silver bromide which is used to make photographic films.

Silver nitrate compound is used in making antiseptics.

Potassium nitrate is used to make gun powder.

Ammonium nitrate is also used in making explosives and blasting agents which are used in mines and quarries. Potassium and sodium nitrate are used as food preservatives.

Chlorides.

Metal chlorides are metal salts containing chloride ions. Such salts include; sodium chloride, potassium chloride and magnesium chloride.

Preparation of metal chlorides by direct methods

Chlorine combines with metals to form metal chlorides for example chlorine combines with sodium to form sodium chloride and with zinc to form zinc chloride.

A lot of heat energy is required for the reactions to take place. The energy is provided by burning the metal.

Reacting a dilute hydrochloric acid and an alkali

Hydrochloric acid + Sodium hydroxide→ Sodium chloride + Water

HC1(aq) + NaOH(aq)→NaCl(aq) + H20(1)

This is a neutralisation reaction.

Reacting a metal and dilute hydrochloric acid in which hydrogen gas is evolved

Zinc + Hydrochloric acid→Zinc chloride + Hydrogen gas

Zn(s) + HC1(aq)→ ZnC12(aq) + H2(g)

Reacting a carbonate and a dilute hydrochloric acid in which carbon(IV) oxide gas is evolved

Lead(II) carbonate + Hydrochloric acid →Lead(II) chloride + Water +Carbon(IV) oxide

PbCO3(s) + 2HC1(aq)→PbC12(s) + H2O + CO2

Insoluble metal chloride can be prepared by reacting dilute hydrochloric acid with an aqueous salt of lead or silver e.g

Hydrochloric acid + Lead(II) nitrate→Lead(II) chloride + Nitric(IV) acid

2HC1(aq) + Pb(NO3)2(aq) →PbC12(s) + 2HNO3(aq)

Hydrochloric acid + Silver nitrate→Silverchloride + Nitric(IV) acid

HC1(aq) + AgNO3(aq)→AgC12 +HNO3(aq)

This method of preparing insoluble metal chloride is called precipitation or double decomposition. It involves two ionic solutions exchanging their anions resulting in the formation of a precipitate, an insoluble compound.

Reacting an oxide and dilute Hydrochloric acid in which salt and water are formed hence it is a neutralisation reaction

Copper(II) oxide + Hydrochloric acidCopper(II) chloride + Water

CuO(s) + 2HC1(aq)→ CuC12(aq) + H20(1)

When dry chlorine is passed over a heated iron, black crystals are observed. The black crystals are deposited in the on the sides of the combustion tube. The black crystals sublimes as brown fumes when the end of the tube is warmed.

Iron(III)chlorine is formed when chlorine is passed over hot iron

Iron + Chlorine + Iron(III) chloride

2Fe(s) + 3C12(g)→ 2FeC13(s)

(black crystals)

Iron (III)chloride sublimes as brown fumes and are collected in the receiver. The whole apparatus should be dry because iron (III) chloride reacts with water vapor in the atmosphere to form hydrogen chloride gas which dissolves in water vapour to form hydrochloric acid

When dilute hydrochloric acid is added to silver nitrate solution, a white crystalline solid is observed

Silver Nitrate + Hydrochloric Acid→Silver Chloride + Nitric(V) Acid

AgNO3(aq) + HC1(aq)→ AgCl(s) +HNO3(aq)

Chemical properties of metal chlorides

Effect of heat on metal chlorides

Common metal chlorides include sodium chloride, magnesium chloride, aluminum chloride, silicon (IV) chloride.

Aluminum chloride exists in some cases as a dimer Al2C16 sodium chloride and magnesium chloride are solids with high melting and boiling points because of the large amount of heat needed to break the strong ionic attractions.

When hydrated magnesium chloride is heated, solid magnesium chloride and water are formed.

Hydrated magnesium chloride

Magnesium chloride + Water

MgC12.6H20(s)

MgC12(s) + 6H20(1)

Electrical conductivity

Sodium and magnesium chloride are ionic compounds and hence undergo electrolysis when they are molten. Electricity is carried by the movement of the ions, and their discharge at the electrodes.

Uses of metallic chlorides

Sodium chloride is used to add taste to food.

Ammonium chloride is used in the manufacture of dry batteries, common salt.

Aluminum chloride is used in petroleum industry

Calcium chloride is used in the extraction of sodium metal

Magnesium chloride is used for dust control and road stabilization

When magnesium chloride is mixed with hydrated magnesium oxide, a hard material (sorel cement) is formed. This is used in fire extinguishers.

Sulphates

Sulphates are salts of sulphuric(VI) acid that contains the anion (S042-). They are formed from the reaction of sulphuric(VI) acid with a metal to form a metal sulphate.

Preparation of sulphates

Soluble sulphates can be prepared by

a) Action of dilute sulphuric(VI) acid on a metal

Zinc + Sulphuric(VI) acid

Zinc sulphate + Hydrogen gas

Zn(s) + H2SO4(aq)

ZnSO4(aq) + 112(g)

When zinc granules are added to dilute sulphuric(VI) acid, an effervescence is observed. Effervescence stops when zinc is added in excess.

When filtrate is evaporated and left to cool large crystals are observed.

Zinc reacts with dilute sulphuric(VI) acid, to form zinc sulphate and hydrogen.

Effervescence stops on adding excess zinc to the acid, because all the acid has been used up.

When excess water is evaporated and the solution left to cool slowly, crystals are formed.

b) Action of dilute sulphuric(VI) acid on a hydroxide

Zinc hydroxide + Sulphuric(VI) acid→Zincsulphate + Water

Zn(OH) 2(aq) + H2SO4(aq)→ZnSO4(aq) + 2H20(1)

Action of dilute sulphuric(VI) acid on an oxide

Sulphuric(VI) acid + Copper(II) oxide→Copper(II) sulphate + Water

H2SO4(aq) + CuO(s)→CuSO4(g) + H20(1)

Action of dilute sulphuric(VI) acid on carbon(IV) oxide gas produced

Zinc carbonate + Sulphuric(VI) acid →Zinc sulphate + Water + Carbon(IV) oxide

ZnCO3(s) + H2SO4(aq)→ZnSO4(aq) + H2O + CO2(g)

Insoluble sulphates are prepared by adding sulphuric(VI) acid in acqueous lead or barium ions for example additions of sulphuric(VI) acid to barium chloride solution.

When dilute sulphuric(VI) acid is added to lead(II) nitrate, a white precipitate is observed. The white precipitate is insoluble.

It is formed as a result of hydrogen ions and lead(II) ions exchanging their anions in the solution

Lead(II) nitrate + Sulphuric(VI) acid →Lead sulphate + Nitric(V) acid

Pb(NO3) 2(aq) + H2SO4(aq) →PbSO4(s) + 2HNO3(aq)

Chemical properties of sulphates

Effect of heat on sulphates

Iron(II)sulphate decompose on heating to produce sulphur(IV) oxide and sulphur(VI) oxide gases. Iron(III) oxide residue is formed.

Iron(II) sulphate→ Iron(III) oxide +Sulphur(IV) oxide + Sulphur(VI) oxide

2FeS 04(s)→Fe203(s) + S02(g) + S03(g)

Hydrated copper(II) sulphate lose water of crystallization on heating to form anhydrous copper(II)sulphate.

Hydrated copper(II) sulphate→Anhydrous copper(II) sulphate + Water

CuS0 4.5H20(s) → CuSO4(s) +5H2O.

On strong heating, the anhydrous copper II Sulphate decomposes to copper II Oxide and sulphur trioxide.

CUSO4(s)→CuO(s) + SO3(g)

The gas turns wet litmus paper red.

Zinc sulphate decomposes2 on the same way on heating.

ZnSO4(s) →ZnO(s) + SO3(g)

When hydrated iron II Sulphate is gently heated, it gives up its water of crystallization to give anhydrous iron II Sulphate.

FeSO4.7H2O(s) → FeSO4(s) + 7 HO (l).

(Light green) (white)

The colour of the iron II Sulphate changes from light green to white.

On strong heating, the unhydrous iron II sulphate decomposes to form a reddish brown iron oxide, sulphur trioxide and sulphur dioxide.

2FeSO4(s) → Fe2O3(s) + SO3(g) + SO2(g)

The sulphur dioxide can be identified by passing it over filter paper dipped in acidified potassium dichromate paper, where it will turn into green colour.

When hydrated iron III sulphate is gently heated, it releases its water of crystallization to form a reddish brown solid which is unhydrous iron III Sulphate.

Fe(SO4)3.9H2O(S) → Fe2(SO4)3(s) + 9H2O(l)

On strong heating, the anhydrous iron III Sulphate decomposes to a reddish – brown iron III Oxide and sulphur trioxide;

Fe2(SO4)3(s)→ Fe2O3(s) + 3SO3(g). Sulphur trioxide turns wet blue litmus paper red.

Test for sulphates.

Barium ions and lead ions are used to test for the presence of sulphate ions because they form insoluble sulphates

When testing for a sulphate, a dilute HCl acid is added followed by barium chloride. The formation of a precipitate confirms that the test solution is a sulphate.

Ba2+ (aq) + SO42- (aq) → BaSO4(s).

The reason why we add the acid is to eliminate any sulphite and carbonate which may be present and can also form a precipitate.

Ba2+(aq) + SO32-(aq) → BaSO3(s)

Ba2+(aq) + CO32-(aq) →BaCO3(s).

The acids reacts with any sulphites and carbonates, which if present would also forms precipitates with barium ions and thus eliminates them.

Uses of metal sulphates.

Calcium sulphate is used as plaster of paris to make moulds, plaster, give accurate production of shapes, etc.

Iron II Sulphateis used to make tablets for patients with iron deficiency redoxide as a pigment

To make potassium hexacyanoferrate which reacts with iron III ions to form a dark blue insoluble solid called Prussian blue which is used extensively as a dye for blue print paper ink

As a weed killer

For treating sewage and water

To coagulate blood in slaughter houses

In turning leather

As fungicides

Barium sulphate is used in the following ways.

As pigment in white paints

Taking X- Ray pictures in digestive tract

Aluminium sulphate used in;

As a mordant in dyeing

As a size in paper making

Potassium aluminium sulphate

Alum is used in dyeing industry

Used as a coagulant in water treatment

Alum is also used in turning of leather

Hydrated sodium sulphate commonly called Glaubers salt is used as a mild laxative

Copper II sulphate is used as a catalyst in the precipitation of ethanol

Hydrated zinc sulphate (ZnSO4.7H2O) is used in the dyeing industry as an antiseptic and in preserving wood. It is used in zinc plating by electrolysis.

COMPOUNDS OF METAL CHAPTER QUESTIONS

Match the items in list A with their corresponding in List B by writing the correct letter beside the item number;

LIST A

LIST B

Its nitrate is used in manufacture antifungal creams.

Its carbonate if used in removal of water hardness.

Its nitrate is used as fertilizers.

Its hydroxide is used to make anti acid tablets.

Brown ring test.

Its nitrates form black residue when heated.

Its hydroxide is used as fuming agent in fire extinguisher.

Basic oxide.

Acidic oxide

Yellowish solid produced when sodium burn in excess air.

Ammonium radical

Sulphur

Carbon

Copper

Sodium oxide

Sodium peroxide

Silver

Gold

Aluminium

Phosphorous pent oxide.

Aluminium oxide.

Dinitrogen monoxide.

Magnesium

Argon

Germanium

Nitroso Iron (II) sulphate.

Potassium zincate.

Test of carbonate in the laboratory.

(a) (i) What is metal oxide?

(ii) Arrange the following metals in the order of increasing their reactivity; zinc,

magnesium, calcium, copper and mercury.

(b) (i) Which one of the metals in a(ii) above react with steam to form an oxide which is

white when cold and yellow when hot?

(ii) Predict the products formed when oxides of zinc and Magnesium reacts with dilute

sulphuric acid.

Differentiate salts behave differently when heated. Complete the following equations that show the actions of heat on certain salts;

bCO3(s)

Ca (NO3)2(s)

(

NH4)2 CO3(s) 2NH3(g)+

H4Cl(s)

What do you understand by the following terms:

Double decomposition

Neutralization

Basic oxide

Amphoteric oxide

State three commercial uses of calcium oxide

Sodium hydroxide can be prepared directly adding sodium metal to cold water. Satae and explain the observation.

Differentiate between basic hydroxide

(a) What is ionic equation

(b) Write balanced ionic equations of the following chemical reactions

Sodium metal is dissolved in water

Dilute sulphuric and is reacted with solid zinc carbonate

Copper ions are precipitated from solution by carbonate ions

Silver nitrate solution gives a white precipitate when reacted with dilute hydrochloric acid

With the aid of a well balance chemical equations explain what would happen in each of the following reactions

Potassium carbonate is strong heated

Concentrated sulphuric acid added to dry salt and heated

Sodium nitrate heated

Concentrated sulphuric acid slowly acted to the sugar

Soluble alkali added to the soluble salts

Carbon dioxide passed through lime water

Water is added to a white copper (11) sulphate

A glowing splint of wood is lowered into a jar of (i) Hydrogen gas (ii) Carbon dioxide gas

Ammonium chloride heated

Dilute nitric acid was heated with (i) warm copper oxide (ii) Zinc carbonate

With the aid of an equation in each case, explain what will be observed when:-

Carbondioxide gas is passed through lime water.

Lead (II) Nitrate is heated

A piece of sodium metal is dropped in water.

You are given the following symbols of metals’ Zn, Na, Cu, Ag, Mg

State the metal in each case;

Which reacts vigorously with cold water?

Which reacts strongly with steam but not with cold water?

The metal whose carbonate doesn’t decompose on heating.

The metal whose nitrate decomposes leaving a metallic residue.

Write equation for reactions in (a) (i) and (ii).

Arrange the above metals in order of increasing activity.

State and explain the effect of passing excess carbon (IV) Oxide through lime water.

State two uses of sodium hydroxide

Complete and balance the following equations

KNO3 (s)

− + −

Ca(NO3)2(s)

− + NO2(g) + −

AgNO3

Ag (s) + − + −

The insoluble silver chloride can be prepared by the reaction of dilute hydrochloric acid and silver nitrate solution. Write the ionic equation for this reaction.

Match the items in List A with their correspondence responses in List B.

LIST A

LIST B

It nitrate is used in manufacture of antifungal cream.

Its carbonate is used in removing water hardness.

Its nitrates is used as fertilizers.

Its hydroxide is used to make ant - acid tablets.

Brown ring test.

Its nitrates form black residue when heated.

Its hydroxide is used as fuming agent in fire extinguisher.

Basic oxide

Acidic oxide

Yellowish solid produced when sodium burn in excess air.

Ammonium radical

Sulphur

Carbon

Copper

Sodium oxide

Sodium peroxide

Silver

Sodium

Gold

Aluminium

Phosphorous pent oxide

Aluminum oxide

Dinitrogen monoxide

Magnesium

Germanium

Iron (ii) sulphate and sulphuric acid.

Argon

Potassium zincate

Test of carbonate in the laboratory

Given the elements, Aluminium, silver, copper, iron, potassium and zinc;

What is reactivity series.

Arrange these metals in the order of increasing their reactivity.

With the help of chemical equations, show the nitrates of elements of silver and potassium decompose when heated.

Explain how hot concentrated sulphuric acid reacts copper metal?

From the metals in the list above identify metals which do not react with dilute acids.

TOPICAL QUESTIONS ON COMPOUNDS OF METALS.

FORM THREE CHEMISTRY.

SECTION A. 20 MARKS.

MULTIPLE CHOICE QUESTIONS

Which of the following properties generally increases down the group?

Ionization energy

Atomic sixe

Electro negativity

Sodium and zinc

When does a chemist fail to identify a compound between sulphur and

iron?

a black solid is formed

heat is used to join them up

yellow color of sulphur and silvery shinny

the resulting mass is greater than the individual mass of the elements

any of the above A-D does not take place.

Which of the following sets of elements is arranged in order of increasing electro negativity.

A. Chlorine, fluorine, nitrogen, oxygen, carbon

B. Fluorine, chlorine, oxygen, nitrogen, carbon

C. Carbon, nitrogen, oxygen, chlorine, fluorine

D. Nitrogen, oxygen, carbon, fluorine, chlorine

E. Fluorine, nitrogen, oxygen, chlorine, carbon

(iv) Which method could be used to separate the products in the following equation?

Pb(NO33)2(aq) + 2KI(aq) → Pb12(s) + 2KNO3(aq)

Colourless colourless yellow colourless

A. Chromatography B. Crystallisation

C. Distillation D. Filtration E. Condensation

(v) The metal nitrate which will NOT give a metal oxide on heating is

calcium nitrate

silver nitrate

lead nitrate

copper nitrate

zinc nitrate

(vi) Which of the following pairs of compounds can be used in preparation of calcium sulphate?

Calcium carbonate and sodium sulphate

Calcium chloride and ammonium sulphate

Calcium hydroxide and barium sulphate

Calcium nitrate and lead (II) sulphate

Calcium chloride and barium sulphate

(vii) The ionic equation when aqueous ammonium chloride reacts with sodium hydroxide solution is represented as:

NH4(aq) +0H-(aq) +0H (aq) → NH3(g) + H20(I)

Na+(aq) + Cl-(aq) → NaCl (s)

H+(aq) + OH(aq) → H2O(l)

2NH4(aq) + 2Cl-(aq) → 2NH3(g) + 2HCl(g)

X, Y, Z

Y, Z, Y

Z, X, Y

Y, X, Z

X, Z, Y

(ix) The only metal which does not react with dilute hydrochloric acid is:-

Magnesium

Aluminium

Copper

Zinc

Sodium

(x) Brown ring test is a technique used to test the presence of;

A. Sulphate

B. Nitrate

C. Chloride

D. Potassium salts

Match the items in List A with their correspondence responses in List B.

LIST A

LIST B

It nitrate is used in manufacture of antifungal cream.

Its carbonate is used in removing water hardness.

Its nitrates are used as fertilizers.

Its hydroxide is used to make ant - acid tablets.

Brown ring test.

Its nitrates form black residue when heated.

Its hydroxide is used as fuming agent in fire extinguisher.

Basic oxide

Acidic oxide

Yellowish solid produced when sodium burn in excess air.

Ammonium radical

Sulphur

Carbon

Copper

Sodium oxide

Sodium peroxide

Silver

Sodium

Gold

Aluminium

Phosphorous pent oxide

Aluminum oxide

Dinitrogen monoxide

Magnesium

Germanium

Iron (ii) sulphate and sulphuric acid.

Argon

Potassium zincate

Test of carbonate in the laboratory

SECTION B

A solution of chlorine in water is acidic.

Yellow phosphorus is stored under water.

Give an account of the following:

Anhydrous copper (II) sulphate becomes colored when exposed to the air for a long time.

Carbon dioxide can be collected by the downward delivery method

Concentrated sulphuric acid is not used for drying hydrogen sulphide gas.

Sodium metal is kept in paraffin oil.

4. (a) With the aid of chemical equations, explain what will happen when aluminum chloride reacts with water.

5. (a) Asubuhi Njemas child was sick. When she took her to the hospital, she was prescribed some medicine including a bottle of syrup. The bottle was written: shake before you use. What does this statement signify?