ORDINARY LEVEL SECONDARY EDUCATION - Lapeer …207.73.181.231/turrill/choggard/8th Grade/STUDENT... · · 2014-01-22ORDINARY LEVEL SECONDARY EDUCATION ... Manipulation of the results

ORDINARY LEVEL SECONDARY EDUCATION

STUDENT PRACTICALS WORKSHEETS

THE UNITED REPUBLIC OF TANZANIA

MINISTRY OF EDUCATION AND VOCATIONAL TARINING

Introduction:

This workbook contains practical worksheets for students on how to carry out practical activities

prescribed by Tanzania Institute of Education in Biology, Chemistry and Physics ordinary level

syllabi.

Practical work is an excellent and effective way of learning and reinforcing theoretical concepts

in science. However, it must be remembered that practical work can be potentially hazardous

and students must always be aware of this.

Each activity follows this layout:

Aim: The objective of the experiment;

Background Information: An introduction which gives student brief explanation to the

practical activity;

Materials: This lists the materials required for the activity;

Procedure: An outline of steps to be followed and how to record data or observations;

Safety Measures: Outlines procedures to reduce hazards;

Analysis and Interpretation: Manipulation of the results obtained and discussion on its

significant meaning;

Conclusion: Students make their conclusion on practical activity;

Questions for Discussion: These are follow-up questions which allow students to

discover new possible knowledge; and

Reflection and Self Assessment: Students state any application for which skills attained

may be applied.

How to use this book

Teachers who make effective use of practical work and experiments often find that students

learn better. Through practical work, teaching is enhanced and becomes more interesting both

for the learner and the teacher.

This is a tool to be used by Biology, Chemistry and Physics Teachers during their lessons which

involve practicals. It is expected to be used effectively to enhance learning of science by doing.

This book consist practical activities for Biology, Chemistry and Physics subjects;

This book will be kept by Teachers in Biology, Chemistry and Physics Departments;

Teachers will produce copies of the practical activity and give to students one day before

the lesson.

Student Report:

Title: Write the title of practical activity

Results: Enter the results of the practical activity obtained as requested from procedure

Analysis: Work out on the results obtained

Conclusion: Conclude the practical activity based on analysis

Questions: Knowledge beyond the practical activity

Reflection: Application of the skills and knowledge attained

Usage of locally available materials:

Science does not have to use expensive or complex resources. It can be taught in the simplest

fashion using empty tins, spirit burners, a few test tubes, plastic drink bottles and materials from

home. For each subject lists of locally available materials can be obtained from Teachers

Practical Guide using locally available materials.

CHEMISTRY WORKSHEET

TABLE OF CONTENTS CHANGES OF STATES OF MATTER

PREPARATION OF A BINARY COMPOUND

PERCENTAGE COMPOSITION OF OXYGEN IN AIR

COMBUSTION OF DIFFERENT SUBSTANCES IN AIR

CONDITIONS FOR RUSTING

EFFICIENCY OF FUELS

CONVERSION OF FORMS OF ENERGY

DISPLACEMENT OF METALS

PRECIPITATION REACTIONS (DOUBLE DECOMPOSITION)

CAUSES OF HARDNESS OF WATER

REACTION OF ACIDS WITH METALS, METAL CARBONATES, METAL OXIDES

AND METAL HYDROXIDES

EFFECT OF HEAT ON SALTS

MEASUREMENT OF MOLAR QUANTITIES OF SUBSTANCES

STANDARDIZATION OF HYDROCHLORIC ACID SOLUTION

STANDARDIZATION OF SODIUM HYDROXIDE SOLUTION

DETERMINATION OF THE WATER OF CRYSTALLIZATION OF A HYDRATED

COMPOUND

DETERMINATION OF THE PERCENTAGE PURITY OF SODIUM HYDROXIDE

CATEGORIES OF ELECTROLYTES

MIGRATION OF IONS TOWARDS OPPOSITE ELECTRODES

PREFERENTIAL DISCHARGE OF IONS AT THE ELECTRODES

FARADAY'S FIRST LAW OF ELECTROLYSIS

FARADAY'S SECOND LAW OF ELECTROLYSIS

ELECTROPLATING OF METALLIC MATERIAL

THE EFFECT OF CONCENTRATION ON THE RATE OF A CHEMICAL

REACTION

THE EFFECT OF TEMPERATURE ON THE RATE OF A CHEMICAL REACTION

THE EFFECT OF A CATALYST ON THE RATE OF A CHEMICAL REACTION

THE EFFECT OF SURFACE AREA ON THE RATE OF A CHEMICAL REACTION

REVERSIBLE CHEMICAL REACTIONS

REACTIVITY SERIES OF METALS

REACTION OF METAL OXIDES WITH ACIDS AND ALKALIS

REACTION OF METALS WITH WATER

REACTION OF SODIUM HYDROXIDE WITH SOLUTIONS OF COPPER (II), IRON

(III), AND ZINC SALTS

PREPARATION OF METAL NITRATES FROM METAL CARBONATES, OXIDES

AND ALKALIS

TREATMENT OF ALKALINE SOIL USING AMMONIUM SULPHATE

EFFECT OF HEAT ON SALTS

ACTION OF DILUTE ACIDS ON SALTS

THE ACTION OF CONCENTRATED SULPHURIC ACID ON A SOLID SALT

SAMPLE

PRECIPITATION OF IONS FROM THEIR SOLUTION BY USING SODIUM

HYDROXIDE OR AMMONIA SOLUTION

CONFIRMATION TESTS FOR CATIONS AND ANIONS

CHANGES OF STATES OF MATTER Aim To investigate how temperature changes affect ice, liquid water, and water vapour.

Background Information Matter exists in three basic states: solid, liquid, and gas. Reversibility of matter allows scientists to convert matter into useful substances in daily life. How can this phenomenon be demonstrated in the laboratory?

Materials Beaker (250 ml), Bunsen burner (or any other heat source), wire gauze, watch glass (or any other cover), tripod stand, ice and stop watch.

Procedure 1. Half-fill a beaker with ice. 2. Put the beaker with the ice on a heat source and heat it for five minutes. Observe what

happens to the ice. 3. Cover the beaker with a watch glass (or any other cover) and continue heating for another

five minutes. 4. Remove the cover from the beaker after five minutes and record what you observe.

Analysis and Interpretation

1. What did you observe within the first five minutes of heating? 2. What did you observe on further heating? 3. What physical processes took place during the experiment?

Conclusion Explain the effect of temperature on the physical states of matter.

Questions for Discussion Explain the changes in the states of matter in terms of product formed, distance between particles and forces between particles.

Reflection and Self Assessment

1. Explain the application of changes of physical states of matter in daily life. 2. Explain how some aquatic animals are able to survive in frozen water. 3. If a bunsen burner is not available in the laboratory, what other heat sources can be used?

PREPARATION OF A BINARY COMPOUND Aim To prepare a binary compound from a reaction of iron and sulphur, and test its properties.

Background Information There are many types of substances in our surroundings. Some of these substances are elements such as iron, sulphur, oxygen, and nitrogen, while some are compounds such as table salt and water. It is interesting to find out what will be formed when two elements such as iron and sulphur are chemically combined and the properties of the reactants and products are tested.

Materials Crucible with lid, spatula, 2 beakers (250 ml), 3 test tubes, bar magnet, heat source, paper, chemical balance, wooden splint, match box, iron filings, sulphur powder and dilute hydrochloric acid.

Procedure 1. Measure about 16 g of sulphur powder and 24 g of iron filings and keep them in separate

beakers. 2. Put a half spatula full of sulphur powder in a small test tube and add dilute hydrochloric

acid. Record any changes. 3. Pour a half spatula full of iron filings into a test tube and add dilute hydrochloric acid. Test

any gas evolved using a burning wooden splint. 4. Mix the remaining sulphur powder and iron filings from step 1 in two separate beakers. 5. Place a small sample of the mixture of iron filings and sulphur on a piece of paper and

move a bar magnet back and forth underneath the paper. 6. Take the remaining mixture and put it in a crucible. Heat the mixture strongly until no more

reaction occurs. Observe and record any changes. 7. Place the residue from heating in a beaker and add dilute hydrochloric acid. Record your

observations. 8. Place a small sample of the residue on a piece of paper and move a bar magnet back and

forth underneath the paper.

Safety Measures This experiment must be performed in open air or in a fume cupboard.


1. Record the results from your observations in tabular form showing the action of dilute hydrochloric on iron, sulphur and the residue.

2. Was there any reaction when hydrochloric acid was added to: a. Sulphur powder? b. Iron filings?

3. What did you observe when a bar magnet was moved under: a. The mixture of sulphur and iron? b. The residue?

4. What colour changes did you observe when the mixture was strongly heated?

Conclusion Are the properties of the reacting elements similar to those of the compound formed after the reaction?

Questions for Discussion

1. Why a gas was evolved when dilute hydrochloric acid was added to the iron filings but not when the acid was added to the sulphur powder?

2. Which element moved when a bar magnet was moved underneath the paper holding the mixture? Which element did not move? Why?

3. The colour of the mixture of iron and sulphur is different from the mixture of the residue. Why?

4. A mixture of iron and sulphur responds to a bar magnet, but their compound does not. Why?

5. What gas was produced when dilute hydrochloric acid was added to the compound?


1. What do you think are two other elements that can produce a binary compound when reacted?

2. How can the process of making a binary compound be useful in industry?

PERCENTAGE COMPOSITION OF OXYGEN IN AIR Aim To investigate the percentage composition of oxygen in air, using a burning candle.

Background Information Air is a mixture of different gases. These gases are nitrogen, oxygen, carbon dioxide and rare gases. It is important to find out the percentage of oxygen in the air we breathe in because it is required in making air tanks for mountain climbers, deep ocean divers and use in hospitals. Breathing and combustion are processes that consume oxygen in the air. This experiment intends to find out the percentage of air consumed by a burning candle.

Materials Water, glass trough, beehive shelf, candle, match box and measuring cylinder.

Procedure 1. Light the candle. 2. Tilt the burning candle to allow the melted candle wax to drop on the beehive shelf just

close to its hole. 3. Fix the burning candle on the molten wax on the beehive shelf. 4. Transfer the beehive shelf with the burning candle into the glass trough. 5. Carefully, without putting off the candle, add water to the glass trough to cover the beehive

shelf up to about 3 cm above the beehive shelf. 6. Invert a 500 ml measuring cylinder over the burning candle as shown in figure (a), and

immediately read the level of water inside the measuring cylinder. Record this level of water as L1 (cm3).

7. Observe what happens to the level of water in the measuring cylinder and the candle flame. 8. Record the final level of water after the candle flame extinguishes as L2 (cm3), see figure (b).

Analysis and Interpretation 1. What volume does the initial level of water, L1, represent? 2. What volume does the final level of water, L2, represent?

3. Find the volume of water in the measuring cylinder after the flame has extinguished. 4. Calculate the percentage air consumed by the burning candle using the formula:

.

Conclusion From this experiment, what is the percentage composition of oxygen in air?


1. Why was it necessary to keep the beehive shelf completely immersed in water? 2. Why did the water rise in the measuring cylinder when the candle was burning? 3. Why did the brightness of the flame continue to diminish until it finally extinguished? 4. Why did water stop rising immediately after the flame extinguished? 5. How does the experimental value of the percentage composition of oxygen in air relate to

that found in chemistry books?


1. Consult chemistry books, or other sources to find out the percentage composition of all components of air.

2. Explain the importance of oxygen in your daily life apart from those mentioned in the background information.

3. Do you think the percentage composition of oxygen in air tanks used in hospitals is the same as that in air? Explain.

COMBUSTION OF DIFFERENT SUBSTANCES IN AIR Aim To determine the products formed when substances burn in air.

Background Information When different substances burn in air, new substances are formed. Ashes, water vapour and gases may be formed depending on the nature of the material burned. The properties of the product formed when substances are burned in air are different from the original substances. What are the characteristics of the products of combustion?

Materials Two test tubes, two rubber stoppers with two holes, beam balance, watch glass, rubber connector, filter pump, glass filter funnel, delivery tubes, marker pen, beaker (250 ml), ice, lime water, anhydrous copper (II) sulphate, kerosene, spirit, charcoal, piece of paper, candle, match box, mortar and pestle, kerosene burner, spirit burner and charcoal burner..

Procedure 1. Label one test tube as A and another test tube as B using a marker pen or any visible mark. 2. In a watch glass measure about 2 g of anhydrous copper (II) sulphate using a beam balance. 3. Transfer the anhydrous copper (II) sulphate to a mortar and grind it with a pestle to form a

powder. 4. Transfer the copper (II) sulphate powder into the test tube labeled A. 5. In the test tube labeled B put about 3 ml of lime water. 6. Set up the apparatus as shown in the figure.

7. Light a candle and start a filter pump while keeping the candle burning. After about 2 to 3 minutes record what happens to test tubes A and B.

8. Repeat step 2 to 7, but replace the candle with the following burning materials: kerosene (in kerosene burner), spirit (in spirit burner), charcoal (on charcoal burner) and a piece of paper (on charcoal burner). Record the observations in a tabular form for each burning substance.

Note that: 1. Air circulation should not be blocked by filter funnel around the heated substances. 2. The delivery tube should be in contact with the anhydrous copper (II) sulphate and lime

water.

Safety Measures Spirit should not be exposed to open flame when not in use because it is highly flammable and evaporates easily.

Analysis and Interpretation 1. What was the colour of lime water before and after burning kerosene, spirit, charcoal, a

piece of paper and a candle? 2. What was the colour of anhydrous copper (II) sulphate before and after burning kerosene,

spirit, charcoal, a piece of paper and a candle?

Conclusion What are the products of burning kerosene, spirit, charcoal, a piece of paper and a candle?

Questions for Discussion 1. Why was ice used in this experiment? 2. What other materials can produce similar effects to charcoal when burned in air? 3. What do you think will happen when kerosene or candle are burned in a limited supply of

air? 4. Why does burning charcoal have a different effect on anhydrous copper (II) sulphate

compared to other burned substances?

Reflection and Self Assessment 1. Do you think burning fire wood (which is used at home) can produce the same products as

burning kerosene in air? Explain. 2. What are the environmental effects caused by gases produced from burning fuels? 3. Why it is not advised to sleep in a room with burning charcoal?

CONDITIONS FOR RUSTING Aim To investigate the conditions necessary for rusting.

Background Information Rusting is a chemical reaction whereby iron forms a reddish brown compound called rust. Rust is a hydrated iron oxide. Rusting only occurs in iron materials. Rusting of iron in industrial and domestic materials can cause economic damage. Structures such as car bodies, bridges, utensils and roofing materials are destroyed by rust. Is it possible to prevent or minimize the rusting process?

Materials 4 conical flasks (250ml), sand paper or steel wool, cotton wool, rubber bungs, heat source, anhydrous calcium chloride, diesel or oil, iron nails, distilled water and measuring cylinder.

Procedure 1. Label four conical flasks A, B, C and D by using masking tape and a marker pen. 2. Clean eight small iron nails using sand paper or steel wool, then put two nails into each of

the four conical flasks. 3. Cover the nails in conical flask A using about 10g of anhydrous calcium chloride and close

the conical flask using a rubber bung. 4. Cover the nails in conical flask B using wet cotton wool and close conical flask with dry

cotton wool. 5. Cover the nails in conical flask C using about 50ml of distilled water and 50ml of oil/diesel. 6. Cover the nails in conical flask D using 50ml of boiled water and close the conical flask with

a rubber bung. 7. Observe and record the changes that occur in the iron nails in conical flasks A, B, C and D

every day for three days.

Safety Measures Make sure the iron nails are well cleaned before you put them in the appropriate conical flasks.

Analysis and Interpretation 1. What happened to the iron nails kept in the conical flasks A, B, C and D after three days? 2. In which flask was both water and oxygen absent? 3. In which flask was both water and oxygen present? 4. In which flasks was there no rust formation? Explain.

Conclusion What are the conditions necessary for iron to rust?


1. Why was it necessary to clean the iron nails with sand paper or steel wool? 2. Why did you use boiled water? 3. Why was wet cotton wool used? 4. Why was anhydrous calcium chloride used?

Reflection and Self Assessment 1. How can you prevent rust in your home utensils? 2. Explain the ways that are used to prevent rusting in large structures like bridges or ships. 3. How does the rusting phenomenon bring about an economic loss at home and in industry?

EFFICIENCY OF FUELS Aim To determine the efficiency of different types of fuels.

Background Information Fuel is a substance that can burn and give out energy in the form of heat or light. We need energy for domestic, transport, and industrial purposes. Fuels are used at home for cooking, warming, and drying. It is important to find out which fuel will provide maximum energy within a short time with minimum waste products. Is liquid fuel more efficient than solid fuel?

Materials Retort stand, kerosene burner, kerosene, water, charcoal, matchbox, charcoal burner, chemical balance, thermometer, measuring cylinder (500 ml) and two beakers (500 ml).

Procedure 1. Weigh 500 g of charcoal using a chemical balance and put it into a charcoal burner. 2. Light the charcoal burner until it is red hot. 3. Measure 400 ml of pure water using a measuring cylinder and pour it into a beaker. Read and

record initial temperature, T1, of water using a thermometer. 4. Place the beaker with its content over the red hot charcoal burner and immediately start a

stop watch. Record the temperature after five minutes as T2. 5. Measure 629 ml of kerosene using a measuring cylinder and pour it in the kerosene burner.

Light the burner and ensure that the flame is blue. 6. Measure another 400 ml of pure water using a measuring cylinder and pour it into another

beaker. Record its initial temperature as T3. 7. Place the beaker with its content over the kerosene burner and immediately start a stop

watch. Record the temperature after five minutes as T4.

Note: The density of kerosene is 0.795 g/cm3.

Safety Measures 1. Keep laboratory windows open to ensure circulation of air (good ventilation). 2. Tightly close the kerosene burner because it can explode.

Analysis and Interpretation Determine the amount of energy in each of the two fuels using the formula H = m×c×∆T Where:

m = mass of water

c = specific heat capacity of water (4.18 kJ kg-1 K-1)

∆T = change in temperature

Density of water = 1 g cm-1

Conclusion

1. Which of the two fuels gave out the larger amount of energy?

2. Which of the two fuels is more efficient? 3. Which of the two fuels is easier to handle?

Questions for Discussion 1. Why is the energy content in the two fuels different? 2. How can we minimize the energy loss during cooking? 3. Could we use petrol or diesel as a liquid fuel in this experiment? Explain.


1. Is this knowledge of fuel choice applicable in real life situations? Explain. 2. Explain the environmental effect of using charcoal as a source of fuel. 3. In which of the fuels did you obtain large amount of waste product? 4. Is food also a fuel? Explain.

CONVERSION OF FORMS OF ENERGY Aim To convert energy from one form to other forms.

Background Information The use of energy cannot be avoided in our daily life. Energy is produced and used at home, at school, in factories and in the bodies of animals and plants for different purposes. There are different forms of energy such as mechanical energy, chemical energy, light energy, heat energy, magnetic energy and sound energy. Chemical energy is stored in materials such as foods, natural fuels (coal, petroleum, and natural gas), batteries and in accumulators. Energy can neither be created nor destroyed, but rather can be changed from one form to another. Energy becomes useful for doing work when one form is changed to another. Which forms of energy are obtained when a given quantity of fuel is burnt in air?

Materials Beaker (250 cm3), thermometer (-10ºC to 110ºC), heat source (burner), 1.5 V electric bulb, connecting wires, switch, ammeter, electric buzzer (or bell), AC mains terminal, distilled water and fuel (for use in the burner).

Procedure 1. Pour 200 cm3 of distilled water into a 250 cm3 beaker. Take the initial temperature of the

water using a thermometer. Heat the water in the beaker using a burner of your choice (Bunsen burner, kerosene burner, spirit burner etc.). Heat the water for five minutes and record its final temperature. Record any forms of energy released by the burner.

2. Connect the following materials in a circuit: one 1.5 V electric bulb, connecting wires, 6V DC source (e.g. DC mains with an adapter, accumulator or batteries), switch and ammeter.

3. Hold the bulb in your finger before and after joining the circuit. Record any forms of energy released by the bulb.

4. Connect an electric buzzer and a switch to AC mains terminals. Complete the circuit and record any energy change.

Safety Measures 1. Do not strike a match if there is any gas leakage in the classroom. 2. Be careful when using a spirit burner because spirit is highly flammable.

Analysis and Interpretation 1. What can be done to energy in order to get the form that is useful for a particular work? 2. What were the energy transformations observed in this experiment?

Conclusion What are the possible advantages of changing energy from one form to another?

Questions for Discussion 1. Energy can neither be created nor destroyed. Explain.

2. Why was it necessary to record the temperature change of the water in this experiment? 3. Was there any energy lost in this experiment? If so, explain.

Reflection and Self Assessment 1. Which part of the procedure was difficult to perform? How can it be improved? 2. How are energy changes used in everyday life? 3. Some machines are less efficient than others in transforming energy. Explain.

DISPLACEMENT OF METALS Aim To investigate the displacement of one metal by another from a compound.

Background Information Displacement is the reaction whereby an element replaces another from a compound. This change usually involves the more reactive metal displacing the less reactive metal. The concept of displacement has applications in the liming process and the storage of solutions in metal containers.

Materials Three beakers (250 ml), sand paper, iron container, copper (II) sulphate solution, silver nitrate solution, zinc foil and copper rod.

Procedure 1. Put 50 ml of copper (II) sulphate solution in an iron container and keep it aside for a week.

Observe and record any changes. 2. Clean a copper rod using sand paper, and place the rod in a beaker containing 50 ml of silver

nitrate solution for a week. Observe and record any changes. 3. Put zinc foil in a 250 ml beaker containing 50 ml of copper (II) sulphate solution and keep it

aside for a week. Observe and record changes.

Safety Measures Avoid getting silver nitrate solution on your skin, clothes, and papers because it can permanently stain.

Analysis and Interpretation From your observations explain what happened when:

1. Copper (II) sulphate solution was put into the iron container? 2. Zinc foil was put into copper (II) sulphate solution? 3. Copper rod was put into silver nitrate solution?

Conclusion From this experiment, which metal is more reactive between the following pairs:

1. Copper and iron? 2. Zinc and copper? 3. Silver and copper?

Questions for Discussion 1. Provide ionic equations for the observations made in each step of the procedure. 2. Briefly explain what would happen if a copper rod is put into a solution of zinc sulphate. 3. Arrange the following metals in the order of increasing reactivity: copper, iron, silver, zinc.

Reflection and Self-assessment 1. Provide some real life applications of displacement reactions. 2. Why is it not advised to store a solution of copper (II) sulphate in an iron container?

PRECIPITATION REACTIONS (DOUBLE DECOMPOSITION) Aim To prepare an insoluble salt by double decomposition or the precipitation method.

Background Information A precipitation reaction is a reaction in which two soluble compounds combine to give an insoluble compound. This type of reaction is called double decomposition because both reactants decompose before the parts reunite to form new products. The reaction is important because it is very rapid and has practical applications in salt formation.

Materials Two test tubes, measuring cylinder, three beakers, silver nitrate solution in a brown reagent bottle, sodium chloride solution, barium chloride solution, copper (II) sulphate solution, and two droppers.

Procedure 1. Put 5 ml of sodium chloride solution in a test tube. To this solution add two drops of silver

nitrate solution. Shake the mixture and allow it to settle. Observe and record what happens in the test tube.

2. Put 5 ml of copper (II) sulphate solution in another test tube. Add four drops of barium chloride solution. Shake the mixture and allow it to settle. Observe and record what happens in the test tube.

Safety Measures Do not allow silver nitrate solution to come into contact with your fingers, cloth, books or papers. They might get permanent stains.

Analysis and Interpretation Write the balanced molecular equation and ionic equation for the reactions of:

1. Silver nitrate solution with sodium chloride solution. 2. Barium chloride solution with copper (II) sulphate solution.

Conclusion What are the insoluble salts that you prepared in this experiment?

Questions for Discussion 1. How could you separate the products in the two experiments? 2. Why did you use the precipitation method to prepare these salts?

Reflection and Self Assessment What is the application of precipitation reactions?

CAUSES OF HARDNESS OF WATER Aim To investigate the causes of hardness of water and how to remove them by boiling and chemical methods.

Background Information Water dissolves many mineral substances. Water is said to be hard if it contains dissolved salts of a particular kind. There are two types of hardness of water: temporary hardness and permanent hardness. Hard water does not produce lather with soap. It is important to investigate substances which cause water to be hard because these may cause wastage of soap during the washing of clothes.

Materials Eight hard glass test tubes (Pyrex), beaker, prepared soap solution, source of heat, test tube holder, filter paper, filter funnel, test tube rack, 0.5 M calcium sulphate solution, 0.5 M magnesium sulphate solution, 0.5 M sodium chloride solution, 0.5 M calcium chloride solution, 0.5 M sodium carbonate solution and 0.5 M calcium hydrogen carbonate solution.

Procedure 1 1. Take five clean test tubes and label them A, B, C, D, and E. 2. Prepare about 50 cm3 of soap solution in a large beaker. 3. Mix 10 cm3 of 0.5 M sodium chloride solution and 10 cm3 of soap in test tube A. Shake the

contents thoroughly. Observe and record any change. 4. Put 10 cm3 of 0.5 M magnesium sulphate solution and 10 cm3 of soap in test tube B. Shake

the contents thoroughly. Observe and record any change. 5. Put 10 cm3 of 0.5 M calcium chloride solution and 10 cm3 of soap in test tube C. Shake the

contents thoroughly. Observe and record any change. 6. Put 10 cm3 of 0.5 M calcium sulphate solution and 10 cm3 of soap in test tube D. Shake the

contents thoroughly. Observe and record any change. 7. Put 10 cm3 of 0.5 M calcium hydrogen carbonate solution and 10 cm3 of soap in test tube E.

Shake the contents thoroughly. Observe and record any change.

Procedure 2

1. Take four test tubes and label them F, G, H, and I. 2. Put 10 cm3 of 0.5 M magnesium sulphate solution in test tube F. Boil the solution for one

minute and leave it to cool. Filter if necessary and add 10 cm3 of soap solution to the filtrate. Shake the contents thoroughly. Observe and record any changes.

3. Put 10 cm3 of 0.5 M calcium sulphate solution in test tube G. Boil the solution for one minute and leave it to cool. Filter if necessary and add 10cm3 of soap solution to the filtrate. Shake the contents thoroughly. Observe and record any changes.

4. Put 10 cm3 of 0.5 M calcium hydrogen carbonate solution in test tube H. Boil the solution for one minute and leave it to cool. Filter if necessary and add 10cm3 of soap solution to the filtrate. Shake the contents thoroughly. Observe and record any changes.

5. Put 10 cm3 of 0.5 M calcium sulphate solution in test tube I. Add 10 cm3 of sodium carbonate solution. Filter if necessary and add 10 cm3 of soap solution to the filtrate. Shake the contents thoroughly. Observe and record any changes.

Safety Measures Do not direct the test tube at your neighbour or yourself when boiling a solution.

Analysis and Interpretation 1. Which test tube contained temporary hard water? 2. Which test tubes contained permanent hard water? 3. Which method was used to remove permanent hardness from water in this experiment? 4. Which method was used to remove temporary hardness from water in this experiment? 5. What type of metallic ions contributed to:

a. Temporary hardness of water? b. Permanent hardness of water?

6. Write the chemical formula of soap scum.

Conclusion

1. Which substances caused hardness of water in this experiment? 2. Write the balanced chemical equations for the process of removing hardness of water in this

experiment.


1. Why did permanent hard water not form lather with soap even after boiling? 2. Why did soap scum form during this experiment? 3. Why did water form a lather with soap after boiling? 4. What are the merits and demerits of hard water?


1. How can you treat hard water for domestic use? 2. What is the effect of scum formation during the washing of clothes with soap? 3. What are the economic losses caused by the use of hard water?

REACTION OF ACIDS WITH METALS, METAL CARBONATES, METAL OXIDES AND METAL HYDROXIDES

Aim To investigate the products formed when acids react with metals, metal carbonates, metal oxides, and metal hydroxides.

Background Information An acid is a chemical substance, which when dissolved in water, produces hydrogen ions (H+

(aq)) as the only positively charged ions. What products are formed when acids react with metals?

Materials Test tubes, test tube rack, measuring cylinder, dropper, wooden splint, match box, calcium metal, copper (II) oxide, calcium carbonate, sodium hydroxide, dilute hydrochloric acid, phenolphthalein indicator, lime water and cork fitted with a delivery tube.

Procedure 1. Arrange four test tubes in a test tube rack and label them T, U, V, and W. 2. Put copper (II) oxide, calcium carbonate, sodium hydroxide, and calcium metal in test tubes

T, U, V and W respectively. 3. Add 10 cm3 of 2 M hydrochloric acid to each test tube. Observe and record any change in

the test tubes and test for any gases evolved. 4. Add a few drops of phenolphthalein indicator (POP) into test tube V. Observe and record

the colour change of the solution.

Analysis and Interpretation 1. What was the gas produced when calcium metal reacted with hydrochloric acid in test tube

W? Justify your answer. 2. What was the gas produced when calcium carbonate reacted with hydrochloric acid in test

tube U? Justify your answer.

Conclusion What type of reaction took place in the test tubes T, U, V, and W? Write the balanced chemical equation for each reaction.


1. Explain what caused the black solid copper (II) oxide in test tube T to change into a pale green solution with the addition of hydrochloric acid.

2. Why is dilute sulphuric acid not suitable to react with calcium carbonate? 3. What other metal oxides, carbonates, and hydroxides can react with dilute acids?


1. What is the scientific significance of the experiments you have performed? 2. Why do egg shells show effervescence when they are dropped in a beaker containing dilute

hydrochloric acid?

EFFECT OF HEAT ON SALTS Aim To identify the products of the thermal decomposition of salts.

Background Information Thermal decomposition is a reaction where a compound breaks down into simpler substances upon heating. Salts produce more than one product when they thermally decompose. The products may include gases and residues. Can these products from thermal decomposition predict the cation and anion present in a salt?

Materials Hard glass test tubes, watch glasses, test tube rack, test tube holder, Bunsen burner, spatula, lime water, wash bottle, delivery tube, distilled water, potassium dichromate paper, match box, red and blue litmus paper, calcium nitrate (Ca(NO3)2), iron (II) sulphate (FeSO4), copper (II) sulphate (CuSO4), lead (II) carbonate (PbCO3), zinc (II) sulphate (ZnSO4) and ammonium chloride (NH4Cl).

Procedure 1. Put a spatula full of calcium sulphate (CaSO4) into a clean, dry hard glass test tube (Pyrex). 2. Light a bunsen burner and adjust it to get a blue flame. 3. Place a tripod stand with wire gauze over the Bunsen flame. 4. Hold the test tube with a test tube holder and strongly heat the sample of calcium sulphate. 5. Test and identify the gas evolved using litmus paper. 6. Observe the residue while it is hot. Then take the test tube out of the flame and observe the

residue while it cools. 7. Repeat the procedure for the remaining salts listed in the materials section. Record all

observations for each salt in tabular form.

Safety Measures

1. Use hard glass test tubes for direct heating on a Bunsen flame. 2. Avoid directing the mouth of the test tube at yourself or other individuals during heating. 3. You are advised to test lead compounds last because their residue is difficult to remove from

test tubes.


1. What was the colour of the residue you observed from each salt? 2. Which gas was produced from each compound after decomposition? 3. Which colour changes helped you to identify the anions and cations present in the

compounds decomposed by heat?

Conclusion How do you relate the gases produced and the colours of the residues with the cations and anions present in the compound?


1. Did you use any other confirmation tests for gases? What were the results of these confirmation tests?

2. Why were you advised to heat the salt samples using hard glass test tubes? 3. Why is it not advised to point the mouth of test tube at yourself or fellow students when

heating the salt?

Reflection and Self Assessment How does ammonium carbonate differ from the other carbonates used in this experiment?

MEASUREMENT OF MOLAR QUANTITIES OF SUBSTANCES Aim To measure the molar quantities of solid substances and the molar volume of gases by using a chemical balance and molar volume box respectively.

Background Information It is easy to measure the quantities of countable and tangible objects using a chemical balance. Similarly, a molar volume box can be used to measure the volume of gases. However, small and invisible objects like atoms, molecules, ions, and electrons cannot be measured directly. These can be measured indirectly as moles. The mole is a quantity equal to the number of atoms contained in 12g of the carbon twelve isotope. We can count and measure other elements in reference to the carbon twelve isotope.

Materials Chemical balance, meter rule, card board, a pair of scissors, masking tape/glue, sodium carbonate, copper turnings, crucible, heat source, and two beakers.

Procedure 1. Accurately measure 64g of copper turnings on the chemical balance and put them into a

beaker. 2. Weigh an empty crucible. Put a sample of about 200g of sodium carbonate in the crucible

and heat it strongly for a few minutes. Record the weight of crucible and sodium carbonate combined. Heat the sodium carbonate again. Continue heating the sample until the weight of the crucible and sodium carbonate remains unchanged.

3. Accurately measure 106g of the sample of sodium carbonate on the chemical balance and put it into a beaker.

4. Make a square box using card board with dimensions of 28.2 cm x 28.2 cm x 28.2 cm.

Analysis and Interpretation 1. Calculate the number of moles in 64 g of copper turnings. 2. Calculate the number of moles in 106 g of sodium carbonate. 3. Find the volume of the air in the molar volume box in dm3.

Conclusion

1. How can you compare the mass of any substance with the number of moles? 2. How can you compare the volume of any gas with the number of moles?


1. What is the mass of 0.5 moles of sodium carbonate in grams? 2. How many atoms are there in 0.5 moles of sodium metal? 3. How can the molar quantities of elements which exist as a liquid at room temperature be

measured?


1. Why is it important to heat sodium carbonate before it is measured on a beam balance? 2. Why is it important to know the number of moles in substances?

STANDARDIZATION OF HYDROCHLORIC ACID SOLUTION Aim To standardize hydrochloric acid solution by using primary standard sodium carbonate solution.

Background Information A solution, whose exact concentration is not known, can be obtained using another solution called a standard solution. A standard solution is one whose exact concentration is known. There are two types of standard solutions: a primary standard solution and a secondary standard solution. A secondary standard solution is prepared by using a primary standard solution. A primary standard solution should be chemically stable and should not change quickly with time. Therefore, primary standard solutions are used to standardize other solutions. It is necessary to learn how to standardize different solutions because this knowledge is applied in pharmaceutical industries, food factories, water treatment systems and agricultural activities.

Materials Beaker (150 cm3), burette (50 cm3), retort stand, pipette (25 cm3 or 20 cm3), conical flask (250 cm3), dropper, hydrochloric acid solution, 0.14 M standard sodium carbonate solution, methyl orange indicator, pipette filler and white tile.

Procedure 1. Pour about 100 cm3 of dilute hydrochloric acid into 150 cm3 beaker. 2. Transfer hydrochloric acid into a clean 50 cm3 burette up to about 0 ml mark and remove all

air bubbles. 3. Clamp the filled burette on a retort stand and record the initial burette reading. 4. Fill 0.14 M sodium carbonate solution into a 25 cm3 or 20 cm3 pipette using a pipette filler

and transfer the solution into a 250 cm3 conical flask. 5. Add 2 drops of methyl orange indicator into the conical flask. Swirl the contents and record

the colour of the solution. 6. Titrate the standard sodium carbonate solution against hydrochloric acid solution until the

colour of the solution in the conical flask changes. Record the final burette reading. 7. Repeat steps 1 to 6 of the procedure to obtain four more titre values.

Safety Measures Clamp the burette on a retort stand gently to avoid breaking it, but tightly enough to avoid it from falling.


1. What was the colour of the solution in the conical flask before titration? 2. What was the colour of the solution in the conical flask at the end point? 3. What is the average titre value? 4. What is the acid to base mole ratio in this titration reaction? 5. Calculate the molarity of the hydrochloric acid solution in mol/dm3.

Conclusion What is the concentration of the given hydrochloric acid solution in mol/dm3 and g/dm3?


1. Why was it necessary to remove air bubbles before titration? 2. In this experiment, why was methyl orange indicator used instead of phenolphthalein

indicator?

Reflection and Self Assessment 1. What difficulties did you encounter in performing this experiment? Explain how you could

overcome them. 2. Can secondary standard solutions be used to standardize other solutions? Explain.

STANDARDIZATION OF SODIUM HYDROXIDE SOLUTION Aim To standardize sodium hydroxide solution by using a standard solution of oxalic acid.

Background Information Sodium hydroxide is a strong base that readily absorbs carbon dioxide and water from the air. For this reason, the mass of sodium hydroxide cannot be measured accurately. Thus, it is difficult to prepare a primary standard solution of sodium hydroxide because the concentration of sodium hydroxide solution changes with time. Oxalic acid is a useful primary standard solution because it is a stable solid whose mass can be measured accurately. In this experiment, standard oxalic acid solution will be used to standardize sodium hydroxide solution.

Materials Beaker (250 cm3), burette (50 cm3), retort stand, pipette (20 or 25 cm3), conical flask (250 cm3), dropper, standard oxalic acid solution, sodium hydroxide solution, phenolphthalein indicator, pipette filler and white tile.

Procedure 1. Use a 250 cm3 beaker to transfer 0.08 M oxalic acid (H2C2O4) into a clean 50 cm3 burette up

to the 0 ml mark. 2. Remove air bubbles from the tip of the burette. 3. Clamp the filled burette on a retort stand and record the initial burette reading in a table. 4. Use a pipette filler to fill sodium hydroxide solution into a 25 cm3 or 20 cm3 pipette and

transfer the solution into a 250 cm3 conical flask. Record the volume of the pipette used. 5. Use a dropper to add two drops of phenolphthalein (POP) indicator into the conical flask.

Mix the content and record the colour of solution. 6. Titrate the sodium hydroxide solution against the standard oxalic acid solution until the

colour of the solution in the conical flask changes. Record the final burette reading in a table.

7. Repeat steps 1 to 6 to obtain four additional titre values.


Analysis and Interpretation 1. What was the colour of the solution in the conical flask before titration? 2. What was the colour of the solution in the conical flask at the end point? 3. What is the average titre value? 4. What is the mole ratio between the acid and base in this titration reaction? 5. Calculate the molarity of the sodium hydroxide solution in mol/dm3.

Conclusion What is the concentration of the sodium hydroxide solution in mol/dm3 and in g/dm3?

Questions for Discussion 1. Why was it necessary to remove air bubbles from the burette before titration? 2. In this experiment, why was phenolphthalein indicator used instead of methyl orange

indicator? 3. What is the function of indicator in this experiment?


1. What new concepts have you learned from this experiment? 2. Explain real life applications of the concepts you have learned in this experiment. 3. What difficulties did you encounter while performing this experiment? Explain how you

solved them.

DETERMINATION OF THE WATER OF CRYSTALLIZATION OF A HYDRATED COMPOUND

Aim To find out the number of molecules of water of crystallization of hydrated sodium carbonate.

Background Information Different hydrated salts have different numbers of molecules of water of crystallization per molecule of salt. The molecules of water of crystallization help to establish the crystalline structure of the salt. How can the number of molecules of water of crystallization of hydrated sodium carbonate be determined?

Materials Burette (50 cm3), pipette (20 or 25 cm3), dropper, conical flask (250 cm3), white tile, retort stand with clamp, filter funnel, two beakers (250 cm3), hydrated sodium carbonate solution, 0.3 M hydrochloric acid solution, methyl orange indicator, dropper, wash bottle, and distilled water.

Procedure 1. Fill a burette with 0.3 M hydrochloric acid up to the 0 ml mark. Remove air bubbles and

record the initial burette reading. 2. Pipette 25 or 20 cm3 of hydrated sodium carbonate solution into a 250 cm3 conical flask.

Record the volume used in the pipette. 3. Add two drops of methyl orange indicator into the conical flask containing hydrated sodium

carbonate solution and observe the colour change. 4. Titrate the hydrochloric acid solution against the sodium carbonate solution in the conical

flask. Note the colour change at the end point. 5. Record the burette reading at the end point. 6. Repeat steps 1 to 5 for three more titre values. 7. Tabulate your results.



1. In the experiment, what is the function of the: a. White tile? b. Methyl orange indicator?

2. What is the average titre value? 3. Write a balanced chemical equation between hydrochloric acid and sodium carbonate

solution. 4. What is the mole ratio between hydrochloric acid to sodium carbonate? 5. If 7.365g of hydrated sodium carbonate were dissolved to make 250 cm3 of solution,

calculate the concentration of hydrated sodium carbonate solution in g/dm3. 6. What is the molarity of sodium carbonate?

7. What is the molar mass of hydrated sodium carbonate?

Conclusion If the chemical formula of hydrated sodium carbonate is Na2CO3·xH2O, what is the number of molecules of water of crystallization?


1. Why is it necessary to remove air bubbles from the burette before titration? 2. Why is it not advised to use more drops of the indicator during titration? 3. Which one is the standard solution, sodium carbonate solution or hydrochloric acid? 4. What other indicator is suitable for this experiment? 5. What might be the source of errors in this experiment? Suggest ways of how to overcome

these errors.


1. How can you apply the knowledge gained in this experiment to tackle other problems outside the classroom?

2. Which part of this experiment was interesting? 3. Is there any part of this experiment you found difficult? How did you solve it?

DETERMINATION OF THE PERCENTAGE PURITY OF SODIUM HYDROXIDE

Aim To find out the percentage purity of a contaminated sample of sodium hydroxide by titration.

Background Information Very few substances are found pure in nature. Many substances are impure, which means they are contaminated with other substances. In most cases we use mixtures of substances, but in some occasions we need to have a pure substance. It is important to know the extent to which a given amount of a substance is contaminated in order to simplify the purification design. Can the titration method be used to determine the percentage purity of a sample of a contaminated sodium hydroxide?

Materials Beaker (250 cm3), burette (50 cm3), retort stand, pipette (25 cm3 or 20 cm3), conical (250 cm3) flask, dropper, standard sulphuric acid solution, sodium hydroxide solution, methyl orange indicator and pipette filler.

Procedure

1. Put 0.1 M sulphuric acid into a clean 50 cm3 burette up to the 0 ml mark. Remove air bubbles, then clamp the burette to a retort stand and record the initial burette reading.

2. Use a pipette filler to pipette 20 or 25 cm3 of impure sodium hydroxide solution and transfer the solution into a 250 cm3 conical flask. Record the volume used in the pipette.

3. Add two drops of methyl orange indicator into the conical flask. Mix the contents and record the colour of solution.

4. Titrate the standard sulphuric acid solution against the impure sodium carbonate solution until the colour of the solution in the conical flask changes. Record the final burette reading in a table.

5. Repeat steps 1 to 4 of the procedure to obtain four more titre values.

Safety Measures 1. Clamp the burette on a retort stand gently to avoid breaking it, but tightly enough to avoid it

from falling. 2. Handle sulphuric acid solution with care because it is corrosive.

Analysis and Interpretation 1. What was the colour of the solution in the conical flask at the end point? 2. What is the average titre value? 3. What is the acid to base mole ratio in this titration reaction? 4. Calculate the molarity of the sodium hydroxide solution in mol/dm3. 5. Calculate the concentration of sodium hydroxide in g/dm3. 6. Calculate the percentage purity of the sodium hydroxide given that 15 g of the impure

sodium hydroxide was dissolved to make 1 dm3 of solution.

Conclusion What is the percentage purity of the impure sample of sodium hydroxide used to make the alkali solution used in this experiment?


1. Suggest a method to standardize an impure sulphuric acid solution. 2. 56.25 g of a sample of impure sodium carbonate was dissolved to make 250 cm3 of

solution. 20 cm3 of the resulting alkali solution required 25 cm3 of 2.000 M nitric acid for complete neutralization. Calculate the mass of impurities in the given sample. (N = 14, C = 12, Na = 23, O = 16, H = 1).


1. Is the titration method of finding the percentage purity applicable to all bases and acids? Why?

2. Explain the real life applications of the concept of purity and impurity of substances.

CATEGORIES OF ELECTROLYTES Aim To identify electrolytes, non-electrolytes, weak electrolytes and strong electrolytes.

Background Information Electrolytes are ionic substances that can conduct electric current when in a molten or aqueous state. These include salts, acids and alkalis. Electrolytes can be classified as weak electrolytes or strong electrolytes. Some other chemical substances do not allow conduction of electricity even if they are in a molten state or solution. These are known as non-electrolytes. How can the categorization of electrolytes, non-electrolytes, weak electrolytes and strong electrolytes best be achieved?

Materials Connecting wires, carbon electrodes, 1.5 V light bulb, 6 V DC source, seven beakers (250 ml), wash bottle, ethanol, distilled water, sodium chloride solution, sugar solution, dilute hydrochloric acid, sodium hydroxide solution, and kerosene.

Procedure 1. Label seven beakers as A, B, C, D, E, F and G. 2. Put about 100 ml of sodium hydroxide solution, ethanol, dilute hydrochloric acid, kerosene,

sugar solution, sodium chloride solution, and pure water into beakers A, B, C, D, E, F and G, respectively.

3. Set up an electrical circuit as seen in the figure.

4. Dip the electrodes into the solution of beaker A. Observe and record what happens to the

electric bulb. Take out the electrodes and rinse them using distilled water from a wash bottle. 5. Repeat step 4 for the remaining solutions and liquids in beakers B, C, D, E, F and G.

Observe and record what happens to the light bulb for each solution.

Safety Measures Electrodes must be cleaned with distilled water and dried before being used.

Analysis and Interpretation In which beakers was the bulb:

a. Very bright? b. Dim or not bright? c. Off?

Conclusion Classify the solutions and liquids in each beaker as a weak electrolyte, strong electrolyte, or non-electrolyte.


1. What was the function of the bulb in this experiment? 2. Why was the bulb off in some of the solutions? 3. What would happen to the bulb if solid sodium chloride was used instead of aqueous

sodium chloride?


1. Which other substances do you think are strong electrolytes, weak electrolytes, or non-electrolytes?

2. Why is it not advised to touch boiling water in an electric heater?

MIGRATION OF IONS TOWARDS OPPOSITE ELECTRODES Aim To demonstrate the movement of ions in an electrolyte.

Background Information Electrolytes are substances that dissociate to form positive and negative ions when in a molten or aqueous state. Positive ions are called cations, while negative ions are called anions. When an electric current is passed through an electrolyte, the cations and the anions are forced to move in opposite directions. This movement is important since it determines the products to be formed at the electrodes in the industrial production of different substances. What substances will move to each electrode during the electrolysis of copper (II) sulphate solution and potassium permanganate solution?

Materials Potassium permanganate crystals, copper (II) sulphate crystals, potassium nitrate solution, filter paper, two microscope slides, two gator clips, connecting wire, switch, 12 V DC source, dropper and microscope.

Procedure 1. Cut a piece of filter paper to about the size of a microscope slide, but its ends slightly longer

than the slide. 2. Put the filter paper on a microscope slide and add potassium nitrate solution drop-wise on

the paper to cover the whole paper. 3. Put a few crystals of potassium permanganate and copper (II) sulphate at the centre of the

filter paper, about 0.5 cm apart from each other. 4. Attach a gator clip to each end of the piece of filter paper. 5. Connect the gator clips to a switch and a 12 V DC source to make a circuit as shown in the

figure. 6. Cover the piece of filter paper with another microscope slide. 7. Connect the switch to complete the circuit.

8. Use a microscope to observe any movement of colour on the filter paper. Record your observations.

Analysis and Interpretation 1. What colour moved to the cathode? Which ions possessed this colour? 2. What colour moved to the anode? Which ions possessed this colour?

Conclusion Is there any evidence that cations and anions move in different directions during electrolysis? Explain.

Questions for Discussion What would happen if water was used instead of potassium nitrate solution?


1. What other method can be used to demonstrate the movement of ions in an electrolyte? 2. What part of this experiment was difficult? Explain how you would improve it.

PREFERENTIAL DISCHARGE OF IONS AT THE ELECTRODES

Aim To determine the products formed when an electric current is passed through various electrolytes.

Background Information An electrolyte is a substance in a molten or aqueous state that can decompose when an electric current is passed through it. Electrolytes have positive and negative ions. When an electric current is passed through an electrolyte solution, the ions move towards each electrode. The ions may react or change when they reach an electrode. What products will be formed when different electrolytes are decomposed by electricity?

Materials Five U-tubes, carbon electrodes, copper electrodes, connecting wires, 12 V DC power source, sand paper, chemical balance, pipette, copper (II) sulphate solution, starch solution, dilute sulphuric acid, sodium hydroxide solution, potassium iodide solution, distilled water, phenolphthalein indicator and universal indicator.

Procedure 1. Label five clean U-tubes as V, W, X, Y and Z. 2. Put copper (II) sulphate solution into tubes V and W. Add two drops of phenolphthalein

(POP) indicator into tube V only. 3. Put dilute sulphuric acid into tube X and add three drops of universal indicator. 4. Put dilute sodium hydroxide solution into tube Y and add three drops of POP indicator. 5. Put dilute potassium iodide solution into tube Z and add three drops of starch solution. 6. Connect each tube V, X, Y and Z to a 12 V DC power source using carbon electrodes and a

switch. 7. Weigh two copper electrodes after cleaning them with sand paper and distilled water. Record

the weight of each copper electrode. 8. Connect the contents of tube W to a 12 V DC power source and switch using the copper

electrodes. Observe any changes in the electrodes and the electrolyte. 9. Test and identify any gas evolved at each electrode. 10. Remove the copper electrodes from the tube and dry them. Weigh them again and record

any changes in mass.

Analysis and Interpretation Record the observations and inferences for each experiment in tabular form.

Conclusion With the aid of equations, write a summary of the products formed in each tube when an electric current is passed through.


1. What caused the colour changes in U-tubes V, X, Y and Z? 2. Why were there no colour changes in test tube W?

3. Explain the reason for the changes in the masses of the electrodes in test tube W. 4. Why were there different products when copper (II) sulphate solution was electrolyzed using

different electrodes? 5. Which of the reactions in the test tubes V, W, X, Y and Z can be used as the basis for the

process of purifying copper? 6. Which conditions determine the nature of the products obtained in electrolysis?


1. What is the industrial use of the principles you have learned in this experiment? 2. What was the most interesting part of this experiment? 3. Which parts of this experiment were difficult for you? How did you cope with the

difficulties? 4. What new ideas or concepts have you learned in this experiment?

FARADAY'S FIRST LAW OF ELECTROLYSIS Aim To determine the relationship between the mass of copper liberated and the quantity of electricity supplied during the electrolysis of copper (II) sulphate solution.

Background Information Substances known as electrolytes may decompose if electricity is passed through them. During electrolysis, either solid or gaseous products are formed at the electrodes. Since the same quantity of electricity passes through the positive and negative electrodes, the mass of the substance corroded or deposited on the electrodes will be the same. What amount of copper will be formed as a result of passing a certain quantity of electricity through copper (II) sulphate solution?

Materials Sand paper, cotton wool, connecting wires, crocodile clips, piece of cardboard, two copper electrodes (copper foils), ammeter (0 – 2.5 A), rheostat, switch, 12 V DC source, distilled water in a wash bottle, copper (II) sulphate solution and beaker (250 cm3).

Procedure 1. Clean two copper foils of the dimension (5 cm x 3 cm) using sand paper and distilled water.

Rinse them with acetone and then hold them suspended in air for about 30 seconds for the foils to dry.

2. Label one of the copper foils as the anode and the other as the cathode. 3. Use an analytical balance to measure the original mass of the anode to the nearest four

significant figures. 4. Using connecting wires and crocodile clips, connect the following in series: ammeter,

rheostat, switch and 12 V DC source (battery or an accumulator). With the switch open, pour about 100 cm3 of 0.1 M copper (II) sulphate solution into a 250 cm3 beaker.

5. Add about three to five drops of dilute sulphuric acid. Fix the electrodes on a piece of cardboard so that they remain separated.

6. Place the electrodes in the electrolyte. 7. Close the switch and immediately start a stop watch. 8. Adjust the rheostat to get a steady current of 0.21 A. 9. After about 15 minutes, switch off the current and stop the stop watch. Read the exact time

on the stop watch in minutes. 10. Remove the electrodes from the electrolyte. Wash the anode carefully with distilled water and

rinse with acetone to dry it. Measure the final mass of the anode. Compare it with the original mass of the anode.

11. Repeat step 3–10 three times. Each time you repeat the steps, change the electrolysis duration to 30, 45 and 60 minutes. Tabulate the results under the following headings: Experiment Number, Mass of anode (Initial and Final (g)), Time (Start and End, (min)) and Current (A).


1. From your tabulated results, complete the following table:

Experiment Current (A)

Time (s) Quantity of electricity (C)

Mass of Cu corroded from the anode (g)

1

2

3

4

2. Plot a graph of mass (g) of copper corroded from the anode (vertical axis) against quantity

of electricity supplied (C) (horizontal axis). Find the slope of this graph. 3. What is the nature of the graph? 4. Compare the value of the slope obtained from the graph with the electrochemical equivalent

of copper found in chemistry books or other data books.

Conclusion How will the mass of a substance liberated or dissolved from an electrode change with increasing quantity of electricity supplied during electrolysis?


1. How did you maintain the current at 0.21 A throughout the experiment? 2. Why was it necessary to maintain a constant current as the experiment was going on? 3. Why was it essential to clean the electrodes thoroughly before passing the current through

them? 4. Some drops of sulphuric acid were added into the electrolyte. How did it make the reaction

go faster without disturbing it?


1. What problems did you face in this experiment? How did you solve them? 2. What new concepts did you learn from this experiment? 3. Which area in this experiment did you find most interesting or least interesting? 4. Which part of this experiment was too difficult for you to carry out? Explain. 5. What are the real life applications of the concepts you have learned in this experiment?

Explain.

FARADAY'S SECOND LAW OF ELECTROLYSIS Aim To investigate the relationship between masses of elements liberated during electrolysis and their chemical equivalents.

Background Information Some metals exist in more than one valence state. For example, copper exists in different compounds as either copper one or copper two (Cu+ and Cu2+). The variable valence of an element affects its chemical equivalent. Ions with higher valence have lower chemical equivalence. If the same quantity of electricity is passed through solutions of different electrolytes, how will the amounts of the elements corroded or deposited on the electrodes be related to the chemical equivalents of the elements?

Materials Connecting wires, crocodile clips, rheostat, ammeter, 12 V DC source (accumulator or a battery of dry cells), iron nail, copper foil, two carbon electrodes, electric switch, 0.1 M copper (II) sulphate solution, 0.1 M iron (III) sulphate solution, two beakers (250 cm3), two pieces of cardboard, sand paper, cotton wool, wash bottle with distilled water, stop watch and analytical balance.

Procedure 1. Using sand paper and distilled water clean a three inch iron nail, a copper foil and two

carbon electrodes. Dry them thoroughly using cotton wool. 2. Measure initial masses of the iron nail and copper foil to four significant figures using a

balance. 3. Label two beakers as A and B using a marker pen or masking tape. 4. Pour enough 0.1 M copper (II) sulphate solution into beaker A until the beaker is half full.

5. Pour enough 0.1 M iron (III) sulphate solution into beaker B until the beaker is half full. 6. Fix the iron nail and a carbon rod on a piece of cardboard so that they are well separated. 7. Fix the copper foil and a carbon rod on another piece of cardboard so that they are well

separated. 8. Use crocodile clips to join an electric switch, a 12V DC source, an ammeter and a rheostat in

series as seen in the figure. 9. Dip the copper foil and carbon electrode into the solution of beaker A. Dip the iron nail and

the other carbon electrode into the solution of beaker B. 10. Close the switch and immediately start a stop watch. 11. Adjust the rheostat to pass a constant current of 0.5A through the circuit for one hour.

Observe and record any colour change in the electrolyte solutions. 12. Stop the reaction by opening the switch. 13. Remove the copper foil and the iron nail from the circuit. Wash them using distilled water,

and dry them using cotton wool. 14. Measure the final masses of the copper foil and iron nail anodes.

Analysis and Interpretation Put all your experimental data in tabular form. Answer the following questions based on your results:

1. What quantity of electricity was supplied in this experiment? 2. What was the change in mass of the anode in:

a. Beaker A? b. Beaker B?

3. What are the chemical equivalents of copper and iron?

Conclusion If the same quantity of electricity is passed through solutions of different electrolytes, what will be the relationship between the masses corroded from the anode or deposited at the cathode and their chemical equivalents?


1. Why do you think it was wise in this experiment to use the corrosion of the anodes instead of deposition of material at the cathodes?

2. What is the difference between the chemical equivalent and the electrochemical equivalent of an element?

3. Why was it necessary to supply the same quantity of electricity through the two electrolytes in this experiment?

4. What would happen if a 6 V DC source was used instead of the 12 V DC source? 5. Considering the charges on the copper and iron ions, which of these elements do you think

will need a greater quantity of electricity to liberate one mole? 6. Write the balanced ionic equations for the reactions which occurred at the anodes in beakers

A and B. What information do these equations provide? 7. If a 1 kg copper anode was immersed in a solution of copper (II) sulphate and was fully

corroded in an electrolysis experiment, what mass of silver immersed in a solution of silver nitrate would also be corroded by the same current?

8. How does the valence of an element help to predict the quantity of electricity needed to produce a certain quantity of the element by electrolytic method?

9. If two cars, A and B, are to be electroplated with silver and chromium respectively, which of these cars will need a greater quantity of electricity?


1. What parts of this experiment were the least or most interesting? 2. What problems did you face in this experiment? How did you solve them? 3. What new concepts did you learn from this experiment? 4. Which part of this experiment was too difficult for you to carry out? 5. What are the real life applications of the concepts you have learned in this experiment?

ELECTROPLATING OF METALLIC MATERIAL Aim To electroplate an iron nail with copper.

Background Information A metal is often plated with another one to protect it from corrosion or to improve its appearance. In the electroplating process, the material to be plated is the cathode and the plating material is the anode. The electrolyte must contain ions of the plating material. When an electric current passes through the electrolyte the plating material is transferred from the anode to cathode. Can electroplating prevent the corrosion of iron?

Materials 9 V DC source, beaker (250 cm3), switch, 0.5 M copper (II) sulphate solution, 2 M sulphuric acid, crocodile clips, connecting wires, iron nails, copper foil, sand paper, and dropper.

Procedure 1. Clean iron nail with sand paper until it shines, wash it with distilled water and dry it using

cotton wool. 2. Clean a copper electrode with sand paper, wash it with distilled water and dry it using cotton

wool. 3. Pour 0.5 M copper sulphate into 250 ml beaker until it is about three quarters full. Add three

drops of 2 M sulphuric acid. 4. Complete the circuit as shown in the figure by connecting the copper electrode to the

positive terminal (anode) and the iron nail to the negative terminal (cathode).

5. Leave the experiment for about 10 minutes and observe the change in the iron nail.

Analysis and Interpretation Using chemical equations show the chemical processes which took place at each electrode.

Conclusion What happened to the iron nail and to the copper electrode?

Questions for Discussion 1. Why were three drops of dilute sulphuric acid added to the electrolyte? 2. What will happen if both electrodes are carbon and the electrolyte is copper (II) sulphate

solution?

Reflection and Self Assessment Explain the application of electroplating in daily life.

THE EFFECT OF CONCENTRATION ON THE RATE OF A CHEMICAL REACTION

Aim To investigate the effect of concentration on the rate of a chemical reaction between sodium thiosulphate and hydrochloric acid.

Background Information It is generally observed that the rate of a chemical reaction depends on the concentration of the reactants. Understanding the relationship between concentration and the rate of chemical reaction is important for the manufacturing industry, where large quantities of chemicals must be produced in short periods of time. How does the change in the concentration of sodium thiosulphate solution affect the rate of its reaction with dilute hydrochloric acid?

Materials Six beakers (100 ml), two measuring cylinders (50 ml), stop watch, white piece of paper, black or blue pen, graph paper, 0.05 M sodium thiosulphate solution, 1 M hydrochloric acid and distilled water.

Procedure 1. Label six beakers as L, M, N, O, P and Q. 2. Pour 0.05 M sodium thiosulphate solution into the six beakers so that 30 cm3, 25 cm3, 20

cm3, 15 cm3, 10 cm3, and 5 cm3 portions go to the beakers, respectively. 3. Pour water into the six beakers so that 0 cm3, 5 cm3, 10 cm3, 15 cm3, 20 cm3, and 25 cm3

portions go to the beakers, respectively. Each beaker should have the same total volume of solution.

4. Draw the letter X on a white paper using a blue or black pen and place it under beaker L. 5. Measure 10 cm3 of 1 M hydrochloric acid using a measuring cylinder. 6. Add the 10 cm3 of 1 M hydrochloric acid into beaker L and at the same time start the stop

watch. 7. Swirl the mixture to mix the contents and look through the solution from above. Stop the

stop watch immediately when the X disappears from sight and record the time it takes for the X to disappear.

8. Repeat steps 4 to 7 of the procedure for the beakers M, N, O, P, and Q.


1. Record the results in tabular form for all six beakers, with a column for the volume of 0.05 M Na2S2O3(aq), V (cm3), the time for X to disappear, t (s), and the rate of reaction, 1/t (s-1).

2. From the results, plot graphs of: a. Volume of 0.05 M Na2S2O3(aq) (V) against time (t); and b. Volume of 0.05 M Na2S2O3(aq) (V) against rate of reaction (1/t).

3. What is the shape of each graph? 4. What does the shape of each graph indicate? 5. What happened to the concentration of sodium thiosulphate after the addition of water? 6. Mention possible sources of error in this experiment.

Conclusion What is the relationship between concentration and the rate of a chemical reaction?


1. Why was the volume of dilute hydrochloric acid kept constant? 2. Write the balanced chemical equation between sodium thiosulphate and hydrochloric acid. 3. What makes the X disappear?


1. Explain situations in daily life where changes in concentration affect rates of chemical reactions.

2. Which problems did you encounter in performing the experiment? What are your suggestions to solve these problems?

THE EFFECT OF TEMPERATURE ON THE RATE OF A CHEMICAL REACTION

Aim To investigate the effect of temperature on the rate of chemical reaction between sodium thiosulphate and hydrochloric acid.

Background Information It is generally observed that the rate of a chemical reaction depends on the temperature of the reaction mixture. The increase in temperature increases the kinetic energy of the reactant particles. This causes the particles to move faster, thus increasing the chances (frequency) of effective collisions between reacting particles. Some reactions such as fermentation, cooking, or some industrial processes depend on the effect of temperature. How do changes in temperature affect the rate of reaction between sodium thiosulphate solution and dilute hydrochloric acid?

Materials Five conical flasks, two measuring cylinders (50 ml), stop watch, thermometer (-10 ºC to 110 ºC), source of heat, white piece of paper, blue or black pen, graph paper, 0.05 M sodium thiosulphate solution, 1 M hydrochloric acid solution and distilled water.

Procedure 1. Label five conical flasks as A, B, C, D and E. 2. Measure 50 ml of 0.05 M sodium thiosulphate (Na2S2O3) solution and pour it into each of

the five conical flasks. 3. Record the temperature of the solution in flask A, which should be at room temperature. 4. Draw the letter X on a white piece of paper using a blue or black pen. 5. Place flask A on the white piece of paper with the X. 6. Measure 10 cm3 of 1 M hydrochloric acid using a measuring cylinder. 7. Add the 10 cm3 of 1 M hydrochloric acid into flask A and at the same time start the stop

watch. 8. Swirl the mixture to mix the contents and look through the solution from above. Stop the

stopwatch immediately when the X disappears from sight and record the time it takes for the X to disappear.

9. Repeat the experiment with the sodium thiosulphate solutions in flask B, C, D, and E by gently heating the solutions to 30ºC, 40ºC, 50ºC and 60ºC respectively. Record all results in tabular form, where you include a column for the temperature of Na2S2O3(aq), T (ºC), the time for the X to disappear, t (s), and the rate of reaction, 1/t (s-1).

Safety Measures 1. Take care when handling concentrated acids. 2. Be careful around open flames during the heating process.


1. From the results, plot graphs of: a. Temperature of Na2S2O3(aq) against time (t); and

b. Temperature of Na2S2O3(aq) against rate of reaction (1/t). 2. What is the shape of each graph? What does the shape of each graph indicate? 3. Mention possible sources of error in this experiment.

Conclusion What is the relationship between temperature and the rate of a chemical reaction?

Questions for Discussion In this experiment, only the sodium thiosulphate solution was heated. What would be the effect of temperature on the rate of reaction if you also heated the hydrochloric acid solution?


1. Explain situations in daily life where changes in temperature affect rates of chemical reactions.

2. Describe the significance of the rate of reaction in food preservation. 3. Which problems did you encounter in performing the experiment? What are your

suggestions to address these problems?

THE EFFECT OF A CATALYST ON THE RATE OF A CHEMICAL REACTION

Aim To investigate the effect of a catalyst on the rate of decomposition of hydrogen peroxide.

Background Information A catalyst is a substance that alters the rate of a chemical reaction but itself remains unchanged at the end of the reaction. Although the catalyst takes part in the chemical reaction, its physical nature may change during the course of a reaction, for example a catalyst may change from coarse particles to fine powder. Catalysts are used in school laboratories as well as in large scale industrial processes. In living things, catalysts are called enzymes. An example of an enzyme is amylase, which is present in saliva. Amylase speeds up the first stage in the breakdown of starch in foods such as bread or potatoes. Some reactions only proceed slowly at room temperature. Can we make these reactions proceed faster by applying a catalyst?

Materials Conical flask with a side arm, two syringes (100 ml), two rubber bungs, sand bath, chemical balance, two retort stands with clamps, stop watch, wash bottle, measuring cylinder, delivery tubes, rubber tubing, filter paper, 20% (v/v) hydrogen peroxide and manganese (IV) oxide.

Procedure 1. Measure 30 ml of 20% (v/v) hydrogen peroxide using a measuring cylinder and put it in a

conical flask. 2. Set up the apparatus shown in the figure.

3. Observe the apparatus for 3 minutes and record what happens. 4. Remove the rubber bung and add 1 g of manganese (IV) oxide. Replace the rubber bung

and immediately start a stop watch. 5. Swirl the flask gently and read the volume of the gas collected in the syringe at 15 second

intervals. Continue observing the syringe until there is no observable change in volume of the gas.

6. Record the volume of oxygen evolved (cm3) using 1g of MnO2 at each 15 second interval of time in tabular form.

7. Weigh a piece of filter paper and record the weight. Use this filter paper to filter the

suspension that is left in the flask. 8. Wash the flask with distilled water and filter the mixture. Repeat this process until all

particles have been removed from the flask. 9. Dry the filter paper together with the residue over a sand bath. Weigh the residue together

with the filter paper. Wait and repeat this process until the mass measured does not change. 10. Repeat the experiment by using 3 g of manganese (IV) oxide. Record the volume of oxygen

evolved (cm3) using 3g of MnO2 at each 15 second interval of time in tabular form.

Analysis and Interpretation 1. Comment on the rate at which the gas was collected when using different masses of

manganese (IV) oxide. 2. What was the final mass of the solid substance collected at the end of the experiment

compared to the mass at the beginning of the experiment? 3. What did you observe regarding the volume of gas before adding manganese (IV) oxide?

Conclusion What is the role of manganese (IV) oxide in the decomposition of hydrogen peroxide and what effect did it have on the rate of reaction?


1. Write a chemical equation representing the experiment. 2. Why did the volume of the gas remain constant at the end of the experiment? 3. Describe any other reaction in which a substance like manganese (IV) oxide can be used.

Reflection and Self Assessment Explain how catalysts are used in human body metabolism.

THE EFFECT OF SURFACE AREA ON THE RATE OF A CHEMICAL REACTION

Aim To investigate the effect of surface area on the rate of chemical reaction between hydrochloric acid and calcium carbonate.

Background Information When all the reactants in a chemical reaction are in the same phase (e.g. gaseous, solution or solid state) the reaction is said to be homogeneous. However, there are cases where the reactants are not in the same phase, and the reaction is said to be heterogeneous. Surface area is only significant in heterogeneous reactions. Reactions involving solids take place on the surface of the solid. Many chemicals used in household supplies, such as soap or baking soda, are sold in powder form instead of large particles because of the effect of surface area on the rate of reaction. Does the surface area of marble chips have an effect on the rate of reaction with hydrochloric acid?

Materials Mortar and pestle, analytical balance, conical flask with a side arm (250 cm3), measuring cylinder (100 cm3), two watch glasses, syringes (100 cm3), stop watch, rubber bung, retort stand with clamp, graph paper, marble chips and 2 M hydrochloric acid.

Procedure 1. Measure 100 cm3 of 2 M hydrochloric acid and put it in a conical flask and set up the

apparatus as shown in the figure.

2. Weigh about 2 g of marble chips in two separate watch glasses. Grind one sample into a fine powder using a mortar and pestle, while the second sample is left as chips.

3. Remove the rubber bung and drop the powdered sample of marble chips into the flask with the acid. Immediately replace the rubber bung and start the stop watch.

4. Record the volume of gas evolved at 10 second intervals for 3 minutes or until the reaction has come to completion. Record your results in tabular form with two rows. One row for the volume of CO2 (cm3) evolved with powdered marble chips and a second row for the volume of CO2 (cm3) evolved with large marble chips. Record the volume in a column for each 10 second interval.

5. Repeat the experiment using the sample of 2 g of large marble chips dropped in a freshly prepared flask of 100 ml of hydrochloric acid.

Safety Measures Take care when handling concentrated acids.


1. From the results obtained, plot a graph of volume of CO2 evolved against time, for both the powdered marble chips and the large marble chips, on the same axes.

2. Comparing the two reactions: a. Which one took less time to go to completion?; and b. Explain the relationship between the slope of the graphs and the reaction rate.

3. From the graph, explain the effect of grinding marble chips into powder on the rate of reaction.

Conclusion What is the relationship between surface area and the rate of a chemical reaction?


1. Eventually no more carbon dioxide gas was evolved, which indicates the reaction has ended. Explain why the reaction eventually came to a stop.

2. In this reaction, we measured the rate at which carbon dioxide gas is evolved. a. Write the balanced chemical reaction between marble chips and hydrochloric acid; and b. Why did we measure the volume of gas instead of measuring the other products?


1. If marble chips are not available, what other substances can you use to do the above experiment?

2. Which problems did you encounter in performing the experiment? What are your suggestions to solve these problems?

REVERSIBLE CHEMICAL REACTIONS Aim To investigate a reaction that could go in both directions.

Background Information There is a group of reactions that can be easily reversed by changing the conditions under which they take place. For example, when ice is heated it melts to form water, and when water is cooled, ice is formed. Such changes are said to be reversible. During the preparation of various important chemical products, one must understand whether the reaction is reversible or irreversible for better control of factors such as quality or product yield.

Materials Spatula, watch glass, mortar and pestle, source of heat, two test tubes, test tube holder, delivery tube, rubber bung, beaker (250 ml), retort stand, cold water and hydrated copper (II) sulphate.

Procedure 1. Take a spatula full of hydrated copper (II) sulphate crystals and grind them to fine crystals

using a mortar and pestle. Observe the colour of the crystals. 2. Put the fine crystals in a test tube and set up the apparatus as shown in the figure. Make sure

the mouth of the test tube with the crystals is slanted downward slightly. Heat the sample gently until there is a colour change. Observe the new colour of the substance.

3. Pour some of the heated sample in a watch glass. Add a few drops of the liquid collected in the second test tube to the sample in the watch glass and observe the colour change.

Analysis and Interpretation 1. What happened when hydrated copper (II) sulphate was heated? What were the products

formed? 2. What happened when a few drops of the collected liquid were added to the residue?

Conclusion What type of reaction is shown in this experiment?


1. What was the remaining residue when the hydrated copper (II) sulphate was heated? Write the chemical equation for this reaction showing the state symbols.

2. Why did you collect the colourless liquid in a test tube immersed in cold water?

Reflection and Self Assessment What are some applications in daily life of the knowledge attained in this experiment?

REACTIVITY SERIES OF METALS Aim To investigate the relative reactivity of metals.

Background Information Most metals react with oxygen, water, and acids to form oxides, hydroxides, and salts. However, metals undergo displacement reactions whereby one metal may displace another metal from its compound. What causes one metal to displace another from its compound? The relative reactivity of metal elements lead to the establishment of an arrangement called the reactivity series. The reactivity series determines the ease with which some metals can be extracted from their oxides by chemical reduction processes. Metals react by a process of electron loss, or oxidation. Metals chemically react as reducing agents. Carbon and hydrogen are actually non-metals, but they are good reducing agents. They are placed with metals in the reactivity series. What is the position of carbon and hydrogen on the reactivity series?

Materials Ten test tubes, two test tube racks, three crucibles, tripod stand, pipe clay triangle, copper turnings, iron (III) chloride solution, copper (II) sulphate solution, iron filings, sodium chloride solution, dilute sulphuric acid, silver nitrate solution, dilute hydrochloric acid, potassium iodide solution, lead (II) nitrate solution, iron (III) oxide, copper (II) oxide, lead (II) oxide, copper (II) nitrate solution, sodium metal, distilled water, aluminum foil, carbon block, magnesium ribbon, calcium granules, candle, match box, blow pipe, heat source, sand paper and phenolphthalein (POP) indicator.

Procedure 1. Your teacher should cut a very small piece of sodium metal and drop it in distilled water

containing phenolphthalein (POP) indicator. Observe the speed of the reaction and the colour changes involved.

2. Label seven clean test tubes as A, B, C, D, E, F and G and place them in a test tube rack. 3. Add the following solutions to the test tubes: distilled water to test tube A; dilute

hydrochloric acid to test tube B; dilute sulphuric acid to test tube C; lead (II) nitrate solution to test tubes D and G; copper (II) sulphate solution to test tube E; zinc chloride solution to test tube F.

4. Clean a small piece of magnesium ribbon and drop it in test tubes A and D. 5. Add a small calcium granule to test tube B. 6. Pour some copper turnings into test tube C using a spatula. 7. Clean a small piece of aluminium foil using sand paper and add it to test tube E. 8. Pour some iron filings in test tube F. 9. Pour copper turnings in test tube G. 10. Place lead (II) oxide powder on a carbon block. Heat the oxide and carbon using a candle

flame strengthened by a blow pipe until only a liquid substance remains. 11. Put a spatula full of iron (III) oxide with another spatula full of copper turnings in a crucible

and mix well. Place the crucible over a pipe clay triangle on a tripod stand and heat very strongly while stirring the mixture. Observe and record all the colour changes in the mixture.

12. Put a spatula full of copper (II) oxide with another spatula full of iron fillings in a crucible and mix well. Place the crucible over a pipe clay triangle on a tripod stand and heat very

strongly while stirring the mixture. Observe and record all the colour changes in the mixture. 13. Label three test tubes as X, Y and Z. 14. Pour about 3 cm3 of silver nitrate in test tube X, and then add 5 cm3 of sodium chloride

solution. 15. Pour about 5 cm3 of lead (II) nitrate solution in test tubes Y and Z. Into test tube Y add

about three drops of iron (III) chloride solution, and into test tube Z add about three drops of potassium iodide solution.


1. Record your experimental data in tabular form showing the experiment done, observations and inferences.

2. What caused the colour change in the water when sodium was dropped in it? 3. In which reactions did you observe colour changes? 4. In which experiments were there no reactions at all? 5. In which experiments did a stronger metal displace a weaker metal from its compound?

Conclusion What evidence can you give to support the fact that some metals are more reactive than others?


1. Why are students not allowed to perform experiment involving the reaction of sodium with water?

2. Carbon can displace some metals from their oxides, even though it is not a metal. Explain. 3. Hydrogen gas can displace copper from its oxide but cannot displace calcium from its oxide.

Explain. 4. Carbon may be used to extract iron from its oxide, but cannot be used to extract aluminium

from its oxide. Explain. 5. Arrange the following elements in the order of decreasing reactivity: lead; zinc; sodium;

potassium; silver; aluminium; calcium; carbon; magnesium; hydrogen. 6. Carbon and hydrogen are non-metals, but are placed on the reactivity series of metals.

Explain. 7. Explain the reason for the colour change which occurred when:

a. Potassium iodide solution was mixed with lead (II) nitrate solution. b. Aluminium foil was placed in copper (II) sulphate solution.


1. Why are most boilers made of copper rather than iron? 2. How does the reactivity series help to choose container material? 3. How would you advise a person who has decided to store sulphuric acid in an iron

container? 4. How is the relative reactivity of elements used in the:

a. Extraction of metals by chemical reduction? b. Manufacture of galvanized corrugated iron sheets? c. Sacrificial protection of iron from rusting?

REACTION OF METAL OXIDES WITH ACIDS AND ALKALIS Aim To examine the amphoteric properties of some metal oxides.

Background Information Metals react with oxygen to form metal oxides. Metal oxides, which react with both acids and bases to form salt and water, are known as amphoteric oxides. On the other hand, metal oxides which react with acid only to form salt and water are referred to as basic oxides. This classification of metal oxides can be done experimentally.

Materials Seven test tubes, spatula, Bunsen burner, two droppers, test tube rack, magnesium oxide, aluminium oxide, zinc oxide, lead (II) oxide, copper (II) oxide, calcium oxide, iron (II) oxide, 1 M hydrochloric acid, and 1 M sodium hydroxide solution.

Procedure 1. Arrange seven test tubes in the test tube rack. 2. Label the test tubes with the chemical formula of the metal oxide (magnesium oxide,

aluminium oxide, zinc oxide, lead (II) oxide, copper (II) oxide, calcium oxide, iron (II) oxide). 3. Put a spatula full of the oxide of each metal in the test tube labeled by its respective name. 4. To each oxide in a test tube, add hydrochloric acid drop wise until the solid sample is

covered by the acid. 5. Shake the test tubes and observe what happens. If there is no change, heat the mixture on

the Bunsen burner flame and observe and record your observation in tabular form. 6. Using the same procedure, repeat the experiment using sodium hydroxide solution instead

of dilute hydrochloric acid.

Safety Measures

1. Test tubes must be thoroughly cleaned after the first experiment before doing the next one to avoid contamination.

2. Do not direct the test tube towards your neighbour or your face while heating the mixture.


1. Which metal oxides reacted with: a. Acid only? b. Both acid and base?

2. With the aid of balanced chemical equations, explain what took place in each experiment. 3. What is the product formed when aluminium oxide reacted with sodium hydroxide?

Conclusion

1. From the experiment, which basic oxides : a. Reacted with sodium hydroxide? b. Did not react with sodium hydroxide?

2. Which metal oxides are amphoteric in this experiment?

Questions for Discussion 1. Why was dilute hydrochloric acid used instead of a concentrated hydrochloric acid? 2. Why did some reactions of metal oxides with hydrochloric acid require heating? 3. Is there any other method that you can use to classify metal oxides? Explain. 4. Mention any other amphoteric oxides apart from those obtained in this experiment.


1. The amphoteric oxides reacted with a basic solution to form a complex compound. What causes this to happen?

2. How can the knowledge of amphoteric oxides be applied to the extraction of aluminium?

REACTION OF METALS WITH WATER Aim To prepare metal hydroxides by the direct method.

Background Information The ability of a metal to react with water depends on its position in the reactivity series. Some metals are very reactive while others are moderate or less reactive. Very reactive metals such as potassium and sodium can react vigorously (explosively) with cold water, forming their hydroxides. Calcium reacts slowly with cold water to form calcium hydroxide. Thus, the hydroxides of potassium, sodium, and calcium can be prepared by the direct method. However, not all hydroxides can be prepared by the direct method.

Materials Beaker (500 cm3), measuring cylinder (500 cm3), test tube, litmus papers (red and blue), wooden splint, match box, water, calcium metal and retort stand.

Procedure 1. Pour 200cm3 of distilled water into a beaker. 2. Put a piece of calcium metal into the beaker of water. Observe and record what happens. 3. Fill a test tube with cold water. Invert it over the calcium metal in the beaker while closing

the mouth of the test tube using a thumb and place it as shown in the figure.

4. Clamp the test tube with the cold water using the retort stand. Make sure the test tube remains upright as shown in the figure.

5. Record your observations. 6. Remove the inverted test tube from the beaker while covering its mouth with a thumb. Test

the gas produced using a burning wooden splint and record your observations.

Safety Measures This reaction is slow at room temperature so you should wait for about one hour to have enough products.

Analysis and Interpretation 1. What was the name of the solution formed in the beaker? Write its chemical formula. 2. Which gas was collected in the test tube during this experiment? 3. What was the test of the gas produced? 4. What was the method of gas collection in this experiment?

Conclusion Write the balanced chemical equation for this reaction. Make sure to write the state of each reactant and product.


1. Why did the water level in the test tube change during the experiment? 2. Is it possible for magnesium to react in cold water and give similar results? Explain. 3. Why did calcium react slowly with cold water in this experiment?


1. How is the concept of reactivity applied in industries? 2. Did you encounter any challenges in this experiment? If yes, explain how you solved them. 3. Other metals in the reactivity series do not react with either cold or hot water. Explain.

REACTION OF SODIUM HYDROXIDE WITH SOLUTIONS OF COPPER (II), IRON (III), AND ZINC SALTS

Aim To prepare and classify metal hydroxides by reacting sodium hydroxides with different salts.

Background Information A hydroxide is any inorganic compound that contains the hydroxide group (OH-). One way to prepare metal hydroxides is by the indirect method, whereby an alkali is reacted with salts, such as chlorides, sulphates and nitrates. Metal hydroxides are classified either on their solubility in water or their reaction with acids and bases. What will be formed when copper (II) sulphate, zinc nitrate, and iron (III) chloride react with sodium hydroxide solution?

Materials Four beakers (100 ml), three measuring cylinders, dropper, test tube rack, three test tubes, distilled water, 0.1 M sodium hydroxide solution, 0.1 M copper (II) sulphate solution, 0.1 M zinc nitrate solution and 0.1 M iron (III) chloride solution.

Procedure 1. Measure 2 ml of copper (II) sulphate solution using a measuring cylinder and add it to a test

tube. 2. Add sodium hydroxide solution to the copper (II) sulphate solution drop wise until it is in

excess. Record your observations. 3. Repeat steps 1 and 2 using zinc nitrate solution and iron (III) chloride solution. Record your

observations for each solution.


1. Present your results in tabular form. 2. Using a balanced chemical equation, write what happens when sodium hydroxide is added to

zinc nitrate gradually until excess. Name the compounds formed.

Conclusion Based on the reaction of each metallic salt with sodium hydroxide, classify the substances formed.


1. Why was sodium hydroxide solution added to the salt solutions drop wise? 2. Differentiate between the compound formed by the reaction of sodium hydroxide solution

with zinc nitrate and the compound formed from the reaction of sodium hydroxide with iron (III) chloride.

3. What will be the products formed if potassium hydroxide solution is used instead of sodium hydroxide solution? Explain.

Reflection and Self Assessment What are the real life applications of the concepts you have learned in this experiment?

PREPARATION OF METAL NITRATES FROM METAL CARBONATES, OXIDES AND ALKALIS

Aim To prepare metal nitrates by reacting nitric acid with metal carbonates, oxides and alkalis.

Background Information Nitrates are obtained from nitric acid through different methods of preparation. These include the reaction of dilute nitric acid with metal carbonates, oxides, and alkalis. Metal nitrates are soluble in water; therefore they cannot be prepared by precipitation.

Materials Five beakers, four evaporating dishes, glass rod, measuring cylinder, tripod stand, source of heat, wire gauze, match box, spatula, copper turnings, nitric acid, lead (II) oxide, calcium carbonate and sodium hydroxide.

Procedure 1. Measure exactly 25 cm3 of 2 M sodium hydroxide solution and pour it in a beaker. Add

exactly 25 cm3 of 2M nitric acid solution. Heat the reaction mixture in an evaporating dish to form a saturated solution and leave it to cool.

2. Put a spatula full of calcium carbonate in a beaker. Add 2 M nitric acid solution to it until all the solid dissolves. Heat the solution formed in the evaporating dish to form a saturated solution and leave it to cool.

3. Put a spatula full of lead (II) oxide in a beaker. Add 2 M nitric acid solution to it until all the solid dissolves. Heat the solution formed in the evaporating dish to form a saturated solution and leave it to cool.

4. Put a spatula full of copper turnings in a beaker. Add 2 M nitric acid solution to it until all the solid dissolves. Heat the solution formed in the evaporating dish to form a saturated solution and leave it to cool.

5. Record the results in tabular form. The table should include columns for experiment, observations, and inferences.

Analysis and Interpretation Dilute nitric acid reacts with the following salts to produce different types of nitrates. Complete and balance the following chemical reactions:

1. HNO3 + CaCO3 → 2. HNO3 + NaOH → 3. HNO3 + PbO → 4. HNO3 + Cu →

Conclusion How can you prepare nitrates in the laboratory? Explain using balanced chemical equations.


1. Why is it important to study the preparation of nitrates?

2. It is not advised to prepare nitrates by using concentrated nitric acid. Explain.

Reflection and Self Assessment Where is the knowledge of the preparation of nitrates applied in daily life?

TREATMENT OF ALKALINE SOIL USING AMMONIUM SULPHATE

Aim To investigate the effect of ammonium sulphate on soil pH.

Background Information Soil pH is very useful tool in order to determine if a soil is suitable for plant growth. Soil pH also determines which nutrients will be absorbed by plants in particular soil. Is it possible to lower soil pH by adding ammonium sulphate salt?

Materials Four conical flasks, soil pH kit, alkaline soil sample, pH indicator paper, ammonium sulphate, filter paper, filter funnel, watch glass, and beam balance.

Procedure 1. Take two clean conical flasks and label them A and B. 2. Measure 50 g of a soil sample and put it into conical flask A. 3. Measure another 50 g of the soil sample and put it into conical flask B. 4. Add 25 cm3 of distilled water into each flask and shake well. 5. Filter the soil sample in flask A and test the pH using pH indicator paper. Compare the color

of the indicator paper against a pH color chart to obtain the pH value of the soil sample. Record your observations.

6. Add 1 g of ammonium sulphate to the sample in flask B and shake well. 7. Filter the mixture in flask B and test the pH using pH indicator paper. Compare the color of

the indicator paper against a pH color chart to obtain the pH value of the mixture. Record your observations.

Analysis and Interpretation How did the soil pH change when ammonium sulphate was added to the soil sample?

Conclusion What is the effect of adding fertilizer such as ammonium sulphate to the soil?


1. Write the ionic equation to explain the effect of ammonium salt on the soil sample. 2. Provide examples of chemical substances that can be used to lower the soil pH.

Reflection and Self Assessment What are the locally available materials that you would advise a small farmer to use in order to reduce alkalinity of the soil?

EFFECT OF HEAT ON SALTS Aim To identify the products of the thermal decomposition of salts.

Background Information Thermal decomposition is a reaction where a compound breaks down into simpler substances upon heating. Some salts produce more than one product when they thermally decompose. The products may include gases and residues. Can these products from thermal decomposition predict the cation and anion present in a salt?

Materials Hard glass test tubes, watch glasses, test tube rack, test tube holder, calcium carbonate (CaCO3), calcium nitrate (Ca(NO3)2), iron (II) carbonate (FeCO3), iron (II) sulphate (FeSO4), copper (II) sulphate (CuSO4), copper (II) carbonate (CuCO3), zinc carbonate (ZnCO3), zinc sulphate (ZnSO4), zinc nitrate (Zn(NO3)2), lead (II) nitrate (Pb(NO3)2), lead (II) carbonate (PbCO3), lead (II) sulphate (PbSO4), Bunsen burner, spatula, lime water, wash bottle, delivery tube, distilled water, potassium dichromate paper, match box, red and blue litmus paper.

Procedure 1. Put a spatula-full of calcium sulphate (CaSO4) into a clean, dry hard glass test tube. 2. Light a Bunsen burner and ensure a blue flame is present. 3. Place a tripod stand with wire gauze over the Bunsen flame. 4. Hold the test tube with a test tube holder and heat the sample of calcium sulphate strongly. 5. Test for any gases evolved. 6. Observe the residue when it is hot. Then take the test tube out of the flame and observe the

residue when it is cold. 7. Repeat the procedure for the remaining salts listed in the materials section. Record all

observations for each salt in tabular form.

Safety Measures

1. Only use hard glass test tubes for direct heating on a Bunsen flame. 2. Avoid directing the mouth of the test tube at yourself or other individuals during heating. 3. You are advised to test lead compounds last because the residue from these salts is difficult

to remove from test tubes.


1. What colour residues did you observe from each salt? 2. What effects did the gases evolved from each salt have on litmus paper?

Conclusion Can you use the colour of a residue and the effect of a gas on litmus paper to confirm a particular salt?

Questions for Discussion 1. Did you use any other confirmation tests for gases? What were the results of these

confirmation tests? 2. Why were you instructed to heat the salt samples using hard glass test tubes? 3. Why were you were advised to avoid directing the mouth of test tube at yourself or your

fellow students when heating the salt?

Reflection and Self Assessment How does ammonium carbonate differ from the other carbonates used in this experiment?

ACTION OF DILUTE ACIDS ON SALTS Aim To determine the products formed when solid metal salts react with dilute hydrochloric acid and sulphuric acid.

Background Information Dilute hydrochloric acid (HCl) and dilute sulphuric acid (H2SO4) have a similar action on solid salts. Thus, the two reagents can be used interchangeably. Some of the reactions of dilute acids with metal salts produce gases. Simple metal salts contain one cation and one anion. Which components of a salt determine the kind of gas produced on action with dilute acids?

Materials Test tubes, spatula, cork fitted with delivery tube, lime water, dilute hydrochloric acid, dilute sulphuric acid, calcium carbonate (CaCO3), copper (II) sulphate (CuSO4), calcium nitrate (Ca(NO3)2), iron (II) carbonate (FeCO3), lead (II) carbonate (PbCO3) ammonium chloride (NH4Cl), sodium hydrogencarbonate (NaHCO3), and zinc chloride (ZnCl2).

Procedure 1. Clean two test tubes using a test tube brush and distilled water. Allow them to dry. 2. Half fill one test tube with lime water. 3. Transfer one spatula full of solid calcium carbonate salt to the second test tube. 4. Add dilute sulphuric acid (H2SO4) to the test tube containing solid calcium carbonate

(CaCO3)). Immediately close the test tube with a stopper fitted with a delivery tube and direct the delivery tube into the test tube containing lime water as shown in the figure. Record your observations.

5. Repeat steps 1 to 4 of the procedure using dilute hydrochloric acid (HCl). 6. Repeat the experiment using the solid salts of copper (II) sulphate (CuSO4), calcium nitrate

(Ca(NO3)2), iron (II) carbonate (FeCO3), lead (II) carbonate (PbCO3), ammonium chloride (NH4Cl), sodium hydrogen carbonate (NaHCO3), and zinc chloride (ZnCl2). Record your observations in tabular form showing experiment, observations, and inferences.

Safety Measures Make sure that distilled water is used to clean the test tubes before and after each experiment in order to avoid contamination.


1. What changes were observed in the lime water? 2. Which salts reacted with dilute acid to produce a gas which changed lime water? 3. What was the possible cause of the observed change in the lime water?

Conclusion What products were formed when each metal salt from the experiment reacted with dilute acids?


1. Why was it important to wash the test tubes using distilled water? 2. Why did the reaction between some metal salts and dilute acid not cause a change in the lime

water? 3. Are there any other reagents or materials that can be used to test the gas evolved in this

experiment? If yes, what are they? 4. With the aid of balanced chemical equations, explain what happened when metal salts

reacted with dilute acids in the experiments you performed.


1. What are the advantages and disadvantages in daily life of the gas that changed lime water? 2. How do you compare the effect of dilute acids and concentrated acids on metal salts? 3. What will happen to the lime water if the gas produced in this experiment passes through it

for a long time?

THE ACTION OF CONCENTRATED SULPHURIC ACID ON A SOLID SALT SAMPLE

Aim To identify the products formed when solid salts react with concentrated sulphuric acid.

Background Information Acids are important reagents which are used to identify anions present in salts. It is important to understand appropriate reactions that can be used to identify anions of different salts. In this experiment, the action of concentrated sulphuric acid on chloride, sulphate, carbonate and bicarbonate salts will be investigated.

Materials Test tube rack, 12 hard glass test tubes, spatula, two droppers, litmus papers (red and blue), concentrated sulphuric acid, test tube holder, lime water, cork fitted with delivery tube, source of heat, ammonium chloride, calcium carbonate, hydrated copper (II) sulphate, iron (II) sulphate, lead (II) nitrate and calcium hydrogen carbonate.

Procedure 1. Put one spatula full of ammonium chloride into a clean, dry test tube. 2. Add 3 drops of concentrated sulphuric acid into the test tube containing the sample and

note the changes. If no changes are observed, gently warm the sample in the test tube. Record your observations.

3. Test for the gas evolved by putting wet litmus papers (red and blue) in the mouth of the test tube. Record your observations.

4. Put one spatula full of calcium carbonate into another clean and dry test tube, and then repeat steps 2 and 3

5. Plug a cork fitted with a delivery tube to the mouth of the test tube containing the test sample.

6. Put about 3 cm3 of lime water into another clean test tube. 7. Insert the delivery tube into the lime water. Record the observations. 8. Put one spatula full of calcium hydrogen carbonate into another clean and dry test tube, and

then repeat steps 2, 3, 5, 6, and 7 of the procedure. 9. Put one spatula full of lead (II) nitrate into another clean and dry test tube, and then repeat

steps 2, 3, 5, 6 and 7 of the procedure. 10. Put one spatula full of lead hydrated copper (II) sulphate into another clean and dry test

tube, and then repeat steps 2, 3, 5, 6 and 7 of the procedure. 11. Put one spatula full of iron (II) sulphate into another clean and dry test tube, and then

repeat steps 2, 3, 5, 6 and 7 of the procedure.

Safety Measures

1. Concentrated sulphuric acid is highly corrosive; handle it with care. 2. These experiments must be performed in a fume cupboard.


1. Using a balanced chemical equation relate the colour change of wet litmus paper with the gaseous products formed when ammonium chloride is reacted with sulphuric acid.

2. What gas evolved when calcium carbonate reacted with concentrated sulphuric acid? Support your answer with a balanced chemical equation, and explain how the gaseous product is related to the colour changes of: a. Wet litmus paper. b. Lime water.

3. Write balanced chemical equations for the reaction of sulphuric acid with each of the following salts: a. Hydrated copper (II) sulphate. b. Iron (II) sulphate. c. Lead (II) nitrate.

Conclusion What gaseous product evolves when concentrated sulphuric acid is reacted with:

1. Chloride salts? 2. Carbonate salts? 3. Hydrogen carbonate salts? 4. Sulphate salts? 5. Nitrate salts?


1. What happens when a hydrated salt is treated with concentrated sulphuric acid? 2. Why is the addition of concentrated sulphuric acid to a solid salt considered a preliminary

qualitative test? 3. Comment on gaseous products formed when calcium carbonate and calcium hydrogen

carbonate were treated with sulphuric acid.


1. Why is it not advised to add water to concentrated sulphuric acid? 2. What are the daily applications of the gases formed when concentrated sulphuric acid is

reacted with carbonates, chlorides and nitrates? 3. Which area in this experiment did you find most interesting or least interesting? 4. Which part of this experiment was difficult for you to carry out? Explain. 5. What are the real life applications of the concepts you have learned in this experiment?

Explain.

PRECIPITATION OF IONS FROM THEIR SOLUTION BY USING SODIUM HYDROXIDE OR AMMONIA SOLUTION

Aim To examine the products of the reaction between lead (II) nitrate with sodium hydroxide solution or ammonia solution.

Background Information Not all cations in a salt sample can be identified by flame colour. However, the cations in their salt solutions can be easily identified by the precipitation reaction of the salt solution using sodium hydroxide solution or ammonia solution. Sodium hydroxide and ammonia react differently with salt solutions. Some salts form precipitates in excess sodium hydroxide or ammonia solutions. Other salts form precipitates that dissolve when excess sodium hydroxide or ammonia solutions are added. These kinds of reactions are applied in the treatment of water, such as the removal of hardness of water, dissolved chlorides and precipitation of heavy metals. It is also applicable in the determination of soil mineral composition and in soil pH management through liming. What kinds of products are formed when a salt solution reacts with sodium hydroxide or ammonia solutions?

Materials Four test tubes, test tube rack, spatula, wash bottle with distilled water, lead (II) nitrate, zinc carbonate, calcium chloride, copper (II) nitrate, iron (II) sulphate, iron (III) chloride, sodium hydroxide solution and ammonia solution.

Procedure 1. Take a spatula full of solid lead (II) nitrate and put it in a clean and dry test tube. Add

distilled water to fill a test tube ¾ with solution. Shake the test tube to dissolve the salt. Note the solubility and keep this test tube as your stock solution.

2. Pour about ¼ of the stock solution into a second test tube then add sodium hydroxide solution drop-wise until in excess. Observe and record any changes.

3. Pour about ¼ full of the stock solution to the third test tube then add ammonia solution drop-wise until in excess. Observe and record the changes.

4. Repeat steps 1 to 3 of the procedure using zinc carbonate, calcium chloride, copper (II) nitrate, iron (II) sulphate and iron (III) chloride. Record your observations in tabular form.

Safety Measure Make sure that test tubes are cleaned and dried before any experiment to avoid contaminations.


1. In which salt solutions did the precipitates disappear when sodium hydroxide solution was added in excess?

2. In which salt solutions did the precipitates disappear when ammonia solution was added in excess?

Conclusion What products were formed when salt solutions reacted with sodium hydroxide and ammonia

solution in this experiment?


1. Why did some of the precipitates disappear in excess sodium hydroxide or ammonia solutions?

2. How can you obtain/collect the precipitates formed from the reactions in this experiment? 3. Write a balanced chemical equation for each of the above reactions using sodium hydroxide

and ammonia solution in this experiment?

Reflection and Self Assessment Apart from the application of precipitation mentioned in the background information, what are other applications of precipitation?

CONFIRMATION TESTS FOR CATIONS AND ANIONS Aim To confirm the presence of particular anions and cations in a solution by the addition of various reagents.

Background Information To determine the nature of a salt solution, you must be able to detect the anions and cations present in the solution. After detecting the presence of a particular ion in a solution, it is desirable to carry out further tests to confirm their presence. What confirmatory tests can be carried out on an unknown salt to confirm the presence of particular anions or cations in their aqueous solution?

Materials Twelve test tubes, dropper, test tube rack, distilled water, test tube holder, dilute nitric acid (HNO3), iron (II) sulphate solution (FeSO4), lead (II) nitrate solution (Pb(NO3)2), calcium hydrogen carbonate solution (Ca(HCO3)2), ammonium chloride solution (NH4Cl), copper (II) sulphate solution (CuSO4), sodium carbonate solution (Na2CO3), iron (III) chloride solution (FeCl3), zinc chloride solution (ZnCl2), sodium hydroxide solution (NaOH), potassium ferricyanide solution (K3Fe(CN)6), potassium ferrocyanide solution (K4Fe(CN)6), potassium iodide solution (KI), magnesium sulphate solution (MgSO4), ammonia solution (NH3), silver nitrate solution (AgNO3), barium chloride solution (BaCl2), dilute hydrochloric acid (HCl), sulphuric acid (H2SO4) and litmus paper (red and blue).

Procedure 1: Confirmation tests for cations 1. Clean six test tubes and label them L, M, N, O, P and Q. 2. To test tubes L, M, N, O, P and Q add 5 cm3 of copper (II) sulphate, iron (II) sulphate, iron

(III) chloride, ammonium chloride, lead (II) nitrate, and zinc chloride solutions respectively. 3. Add ammonia solution to test tube L drop-wise until excess. Observe and record the

changes. 4. Add potassium ferricyanide solution to test tube M drop-wise until excess. Observe and

record the changes. 5. Add potassium ferrocyanide solution to test tube N drop-wise until excess. Observe and

record the changes. 6. Add 5 cm3 of sodium hydroxide to test tube O. Heat the mixture and test any gas evolved

using litmus paper. 7. Add potassium iodide solution to test tube P drop-wise until excess. Observe and record the

changes. 8. Add potassium ferrocyanide solution to test tube Q drop-wise until excess. Observe and

record the changes.

Procedure 2: Confirmation tests for anions 1. Clean five test tubes and label them as G, H, I, J, and K. 2. To the test tubes G, H, I, J, and K add about 5 cm3 of lead (II) nitrate, copper (II) sulphate,

ammonium chloride, sodium carbonate, and calcium hydrogen carbonate solutions respectively.

3. Add freshly prepared iron (II) sulphate solution to test tube G, followed by concentrated

sulphuric acid carefully along the side of the test tube. Do not shake the mixture. Allow the mixture to settle. Observe and record any changes.

4. Add barium chloride solution to test tube H, followed by a few drops of dilute hydrochloric acid. Observe and record any changes.

5. Add a few drops of silver nitrate solution to test tube I, followed by a few drops of dilute nitric acid, and then by ammonia solution. Observe and record any changes.

6. Add magnesium sulphate solution to test tube J. Observe and record any changes. 7. Add magnesium sulphate solution to test tube K and boil the mixture. Observe and record

any changes.

Safety Measures

1. Handle concentrated acids with care because they can burn the skin. 2. Avoid getting silver nitrate solution on your skin, clothes, and papers because it can

permanently stain.

Analysis and Interpretation Present your experimental findings in a table with the headings: Experiment, Observations, and Inferences.

Conclusion For each confirmatory test, name the cation or anion confirmed.


1. Give the reason why many insoluble salts can dissolve in dilute nitric acid. 2. Explain why some cations form a precipitate with some reagents, but the precipitate

dissolves in an excess of the reagent.


1. What problems did you face in doing this experiment? How did you overcome them? 2. What new concepts did you learn in this experiment? 3. How are confirmatory tests used in real life?

Local Materials list Apparatus:

Beaker – Metallic Mug, plastic bottle, glass

Watch glass – Lid, Petroleum jelly bottle lid

Test tube – opened light bulb, syringe with plunger and needle removed and narrow end melted shut

Delivery tube – Infusion tube

Filter funnel – upper part of a plastic bottle

Measuring cylinder – calibrated bottle

Spatula – tea spoon

Gas Jar – Water bottle

Chemicals:

Copper (II) sulphate – Mruturutu

Lime water – Chokaa mixed with water

Calcium sulphate – Gypsum

Magnesium sulphate – Epsom salt

Sodium chloride – Table salt

Sodium hydrogen carbonate – Baking powder

Phenolphthalein indicator – local indicator from flower petals

Sodium hydroxide – Caustic soda

Calcium carbonate –

Acids – Battery acid water, vinegar, citric acid

Manganese IV oxide – black powder from dry cell battery

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ORDINARY LEVEL SECONDARY EDUCATION - Lapeer …207.73.181.231/turrill/choggard/8th Grade/STUDENT... · · 2014-01-22ORDINARY LEVEL SECONDARY EDUCATION ... Manipulation of the results