UNIT 1 : Carbon and its inorganic compounds

  • UNIT 1 : Carbon and its inorganic compounds

    Key unit competency

    After studying this unit, I should be able to relate the properties of carbon and its compounds to its uses and describe how the compounds of carbon are prepared.

    Learning objectives

    By the end of this unit, I should be able to:

    • Name the allotropic forms of carbon and relate its properties to its uses.

    • Explain the properties of carbon and its compounds.

    • Prepare, collect and test for carbon dioxide gas.

    • Prepare, collect and test carbonates of different metals.

    • Explain the impact of carbon compounds on the environment.

    • Explain the carbon cycle.

    • Appreciate the importance of natural resources.

    • Develop self – confidence during presentations and a culture of working cooperation together in groups.

    Mind teaser

    Study the pictures below. Identify the things in the pictures.

    Can you tell from what element the items or some of their components in the pictures are made from?

    Introduction

    The items shown in the pictures, fig 1.1 are all made from carbon element. Carbon is one of the most important elements. Carbon exists both naturally as diamond and graphite and in combined forms in many compounds. Coal, oil, natural gas, limestone and other metal carbonates are all compounds of carbon. Charcoal and the positive electrode in dry cells are all composed of carbon.

    1.1 Definition of allotropy

    Activity 1.1

    Search about meaning of allotropy and present your findings to the class. Use the following questions as a hint.
    1. What is allotropy? Get the meaning from the dictionary, textbooks or the internet.
    2. Which elements exhibit allotropy?
    3. Name two allotropes of carbon that you have identified from this activity.

    The facts.

    Atoms in certain solids often may arrange themselves in different patterns. When this occur it results in different forms of that element called (allotropes). Allotropy is therefore the existence of an element in two or more forms. Elements that exhibit allotropy include carbon, sulphur and phosphorus. Allotropes have same chemical properties but different physical properties.

    1.2 Allotropes of carbon and their physical properties

    In Senior 2 under covalent bonding, you learnt about giant covalent structures,their properties and uses. In this subtopic, you are going to learn how covalent bonding brings about the allotropic forms of carbon.

    Physical properties of carbon allotropes and their uses

    Materials

    Activity 1.2

    Gloves, dry cells, ball-pein hammer and a protective sheet of paper.

    Procedure

    1. Spread the protective sheet of paper on the floor.
    2. Wear gloves and use the ball-pein hammer to crush the dry cell.

    3. Remove the black rod at the centre.
    4. Try breaking it into smaller pieces and feel it between your fingers.

    Study questions

    1. What is the colour of charcoal, soot and the rod in the dry cell?
    2. What is the texture of the charcoal and soot?
    3. How are soot and charcoal formed?
    4. Which allotrope of carbon is used to make the rod in the dry cell?

    Discussion corner!

    1. Answer the study questions above.
    2. In groups, discuss the concept of allotropy; consider graphite in dry cells.
    3. Prepare a report and present it to the rest of the class.

    Self-evaluation Test 1.1

    1. What is allotropy?
    2. Name two elements that exhibit allotropy.
    3. Name the allotropes of elements named in question 2 above.
    4. Complete the table below in relation to charcoal.


    5. ___________ is the allotrope of carbon used in dry cells.
    6. Soot and charcoal are formed when ________________.
    7. What is the role of the rod in a dry cell?

    Crystalline forms of carbon

    Research activity

    1. Using textbooks and from the internet, research on the structure, the physical properties and uses of diamond and graphite.
    2. Write a report of your findings and do a class presentation.
    3. Compare your findings with the ones given below.

    The facts

    There are two main crystalline forms of carbon: diamond and graphite. Non-crystalline forms of carbon also exist. They are called amorphous carbon.

    (a) Diamond

    Diamond occurs naturally in many countries, for example Tanzania, South Africa, India and South America. Thus diamond is a natural resource.

    Structure of diamond

    In diamond, each carbon atom is covalently bonded to four other carbon atoms giving a regular tetrahedron shape as shown in fig 1.2. This results in a rigid closely interlocked three dimensional structure. The atoms form a giant atomic structure. All the four valence electrons in diamond are used in the bonding, which greatly contributes to the physical properties of diamond


    Physical properties of diamond

    • It is the hardest known naturally occurring substance. This is because of the strong covalent bonds that hold the atoms of carbon together
    .• When carefully and specially cut, it is colourless crystalline and transparent with a dazzling brilliant lustre. The lustre is caused by its high refractive index
    .• Diamond has a higher density of 3.5g/cm3 compared to that of graphite (2.25g/cm3).This is because the arrangement of atoms in a diamond crystal allows for more atoms to be packed per given space compared to graphite.
    • It has a high melting and boiling points. It melts at 4200°C. This is because all the carbon atoms are bonded by very strong covalent bonds that require a lot of heat energy to break
    .• Diamond does not conduct electricity. This is because it has no delocalized electrons
    .• Diamond is a good conductor of heat because of its strong covalent bonding.

    Uses of diamond

    Can you remember the uses of diamond you learnt in Senior 2?The physical properties of diamond makes it find important applications in the following areas.
    1. Because of its hardness, it is used for making drill tips used when drilling or cutting metals. It is also used to make glass cutters.
    2. Used in making jewellery due to its beautiful sparkling radiance and lustre.


    (b) Graphite

    Graphite is a dark grey shiny crystalline solid. It is the most abundant allotrope of carbon. Graphite occurs naturally but a superior form of it is made by heating anthracite (a variety of coal) in an electric furnace.

    Structure of graphite

    In graphite, each carbon atom is covalently bonded to three other atoms while the fourth electron is delocalised; i.e it is not attached to any particular atom but belongs to the entire structure. These delocalised electrons are free to move through the graphite structure making graphite a good conductor of electricity. Graphiteexists in hexagonal layers of carbon atoms. The layers are held together by weak Van der Waal’s forces of attraction as shown in fig 1.4. This makes them slide easily over each other.


    Physical properties of graphite

    • Graphite is a soft, black and shiny material with a greasy feel.
    • Graphite easily flakes off.
    • Has relatively high melting and boiling points due to the strong covalent bonds joining its atoms together.
    • It has a density of 2.25g/cm3.
    • Graphite is a good conductor of electricity due to the presence of delocalised electrons in its structure
    .• It is opaque. Numerous parallel layers arranged on top of one another blocks light from penetrating through.

    Uses of graphite

    The properties of graphite determine its uses. Some uses of graphite include:

    1. It is used as electrodes in dry cells and fuel cells. This is because it is a good conductor of electricity.
    2. Graphite is used as a carbon raiser in the production of steel. It gives steel its strengthening characteristics.
    3. It is used in advanced high-friction applications such as car brakes and clutches because of its high thermal and electrical conductivity.
    4. It is used to make pencil ‘leads’ when mixed with clay. This is due to the sliding of its layers which enable it to slide on paper when writing.
    5. At high temperatures, graphite can be used as a lubricant in place of grease or oil. This is because graphite has a high boiling point.
    6. In production of paints and shoe polish, the powdered form of lump graphite is used due to its natural water-repellent property. It is the best choice for giving a protective coating on wood or shoe leather.
    7. It is used to make furnace linings, brushes for electric motors and generators.
    8. It is used as a moderator in nuclear reactors.


    Non-crystalline forms of carbon

    Activity 1.3

    Materials

    Pieces of charcoal, soot, gloves, ball-pein hammer and a sheet of paper.

    Procedure

    1. Spread the protective sheet of paper on a work bench.
    2. Place the pieces of charcoal and soot on the bench.
    3. In pairs, feel the charcoal and soot between your fingers.
    • Observe what happens to your fingers.
    4. Crush the pieces of charcoal into smaller pieces and see what happens.

    Non-crystalline forms of carbon are formed by decomposing other substances like wood using heat.
    Charcoal and soot are examples of amorphous carbon.

    (a) Charcoal

    There are two types of charcoal; wood and animal charcoal. Wood charcoal is obtained when wood is strongly heated in absence of air. This process is called destructive distillation of wood. Remember that large-scale and uncontolled charcoal-making leads to destruction of forests.

    My environment, my life!
    Avoid cutting down trees to make charcoal. Use alternative sources of fuel like biogas that is environmentally friendly.

    Animal charcoal is formed when bones are subjected to destructive distillation. Animal charcoal is only about ten percent carbon with the rest being mainly calcium phosphate.
    Other sources of amorphous carbon include:
    • Burning petroleum products in a limited supply of air to form lamp black.
    • Heated sugar in limited supply of air to form sugar charcoal. This can also be formed by dehydrating sugar using concentrated sulphuric acid.
    • Heating coal in the absence of air to obtain coke.
    • Soot is formed when there is incomplete combustion of fuels. It is found e.g in chimneys of houses and lantern lamp.



    Physical properties of charcoal

    Activity 1.4

    Materials

    Pieces of charcoal, two 100 cm3 beakers, solution of a dye, filter paper, filter funnel, Bunsen burner.

    Procedure

    1. Put about 50 cm3 of the dye solution into a beaker.
    2. Place some pieces of charcoal into the beaker and heat for about ten minutes.
    3. Filter the contents of the beaker and observe any changes in the filtrate.

    Study questions

    1. What changes occurred in the solution after it was heated with charcoal?
    2. Explain the changes that took place.



    I have discovered that...

    The intensity of the dye in solution decreased at the end of the experiment. Particles of the dye stick onto the surface of charcoal.

    The facts

    Charcoal adsorbs particles in a dye reducing the intensity of the colour of the dye. The particles of adsorbed substance stick onto the surface of the charcoal. This property is used in domestic water treatment. Some physical properties of charcoal include:
    • Charcoal is a black porous solid.
    • It is soft and has a low density.
    • It can adsorb large volumes of gases and solids.

    Uses of charcoal

    1. Charcoal is used as a source of fuel mostly in developing countries.
    2. Activated charcoal (charcoal that is finely powdered) is used to remove smelly gases in slaughter houses, gas manufacturing plants, large air conditioning systems and airports. This is because activated charcoal has a large surface area and hence can adsorb large volumes of gases.


    3. Lamp black is used in making black ink, paints, carbon paper and as an ingredient in the rubber tyres.

    Self-evaluation Test 1.2

    1. While cooking, if the bottom of the vessel is getting blackened on the outside,it means______________.
    2. Name one non-crystalline form of carbon that you know.
    3. Match the following statements (i.e A and B) using a line.


    4. Explain the following:
    (a) Diamond is used in making tips of drills while graphite is used to make pencil leads.
    (b) Graphite conducts electricity while diamond does not.
    5. Charcoal is a cheap but expensive source of fuel. Explain.
    6. How does excessive use of charcoal as a source of fuel affect our environment?
    7. Explain why graphite is used as a lubricant whereas diamond is not.

    1.3 Chemical properties of carbon

    a) Reaction with oxygen

    Activity 1.5

    Apparatus and reagents

    Deflagrating spoon, Bunsen burner, gas jar of oxygen,calcium hydroxide solution.

    Procedure

    1. Place a piece of charcoal in a deflagrating spoon and heat it strongly using a Bunsen burner flame until the charcoal glows red-hot.

    2. Lower the red-hot charcoal into a gas jar of oxygen. • What do you observe?

    3. After sometime, pass the product in the jar through a solution of calcium hydroxide and observe what happens.


    Study questions

    1. Why did burning stop after sometime?
    2. What caused the white precipitate when the gas formed in the jar was passed through the calcium hydroxide solution

    Discussion corner!

    Discuss the observations you have made and the study questions in activity 1.5 with a classmate and share your conclusions with the rest of the class.

    I have discovered that...

    Charcoal continues to burn in the gas jar even after heating is stopped. The calcium hydroxide solution turned milky due to the carbon dioxide gas formed when charcoal burnt in the gas jar of oxygen.

    The facts

    Carbon burns in enough supply of oxygen to form carbon dioxide gas. The reaction is exothermic and hence burning continues even after heating has been stopped.
    Carbon + oxygen → carbon dioxide + heat
    C(s) + O2(g) → CO2(g) + heat
    When carbon dioxide was passed through calcium hydroxide, it reacted to form calcium carbonate thus white precipitate is seen. The white precipitate would disappear if carbon dioxide gas is bubbled in excess due to formation of calcium hydrogen carbonate.

    Calcium hydroxide + carbon dioxide  Calcium carbonate + water
    Ca(OH)2 (aq) + CO2(g) CaCO3(s) + H2O(l)

    Calcium carbonate + carbon dioxide + water Calcium hydrogen carbonate
    CaCO3(s) + CO2 (g) + H2O (l) Ca(HCO3)2(aq)

    Carbon reacts with limited oxygen supply to form carbon monoxide.
    Carbon + oxygen carbon monoxide + heat
    2C(s) + O2 (g) 2CO (g) + heat

                                            Health check
    Avoid using charcoal stove in a closed room because the built up of carbon monoxide to toxic levels causes suffocation leading to death!

    b) Reaction with carbon dioxide

    Activity 1.6

    Materials
    Charcoal, source of carbon dioxide, source of heat, combustion tube.

    Procedure
    1. Set up the apparatus as shown below.


    2. Heat the charcoal until it is red-hot.
    3. Pass dry carbon dioxide gas through the heated charcoal.
    4. Ignite the gas coming out through the other side of the combustion tube as shown above.

    Precaution: Avoid inhaling the gas produced. It is poisonous!

    Study questions
    1. What observations did you make in the ignition tube?
    2. Write a balanced chemical equation for the reaction that took place in the combustion tube.

    Discussion corner!

    In pairs, discuss the results of the experiment and the study questions above. Write a report and present to the other class members.

    I have discovered that...
    When carbon dioxide is passed through heated charcoal, the charcoal diminishes. A gas is produced that burns with a blue flame.


    The facts

    Carbon dioxide reacts with heated charcoal to form carbon monoxide gas.
    Carbon monoxide in turn burns with blue flame.
    CO2 (g) + C(s) 2CO (g)

    c) Reaction with iron (III) oxide

    Activity 1.7

    Materials and reagents

    Crucible, iron (III) oxide, source of heat, powdered charcoal, tripod stand, wire gauze.

    Procedure
    1. Mix charcoal powder with iron (III) oxide in a crucible.
    • Note the initial colour of the mixture.
    2. Apply heat on the crucible until it is red-hot.
    • Note the colour of the final product.


    Study questions

    1. Describe the color change from that of the inside mixture and that of the final product. Give reasons for these changes.

    Color of the mixture at the beginning.              Color of the mixture after heating.

    2. Write a chemical equation for the reaction that took place.

    Discussion corner!
    In pairs, discuss the observations made and answers to the study questions above and write a report.

    I have discovered that...
    When the mixture of iron (iii) oxide and powdered charcoal is heated, a red glow is produced. The glow continues spreading throughout the mixture even when heating is stopped. The color of the mixture turns from red-brown before heating to grey after heating.

    The facts

    Carbon in the charcoal reacts with iron (III) oxide to form iron metal and carbon dioxide. This is a redox reaction where carbon reduces iron from its oxide and is itself oxidised to carbon dioxide. The reaction is exothermic and hence continues even after heating is stopped.

    Iron (III) oxide + carbon   Iron + carbon dioxide
    2Fe2O3(s) + 3C(s)      4Fe(s) + 3CO2 (g)
    (Red-brown)                          (Grey)

    Work to Do

    Name other metal oxides reduced by carbon to their respective metals

    d) Reaction of carbon with concentrated sulphuric acid and nitric acid

    Activity 1.8

    Apparatus and reagents

    Two test tubes, concentrated sulphuric acid, cork, delivery tube, retort stand, concentrated nitric acid, powdered charcoal, spatula, bunsen burner, tripod stand, calcium hydroxide solution (lime water).

    Procedure

    1. Put a spatulaful of charcoal powder in a test tube.
    2. Add about 10 cm3 of concentrated sulphuric acid, warm and observe.
    3. Repeat the experiment using concentrated nitric acid instead of sulphuric acid.



    Study questions

    1. What observations did you make on reacting each acid with carbon?
    2. Write balanced chemical equations for the reactions that took place.

                          Discussion corner!
    1. In pairs, discuss the observations made.
    2. Come up with an explanation for your observations. Present to the rest of the class.

    I have discovered that...
    When carbon reacts with concentrated nitric and sulphuric acids, bubbles of a gas are produced. The amount of charcoal powder in the test tube eventually decreases.

    The facts

    Carbon reacts with concentrated sulphuric acid to form carbon dioxide, sulphur dioxide and water. With concentrated nitric acid, it forms carbon dioxide, nitrogen dioxide and water as shown below.

    Carbon + sulphuric acid carbon dioxide + sulphur dioxide + water
    C(s) + 2H2SO4(aq) CO2(g) + 2SO2(g) + 2H2O(l)

    Carbon + nitric acid carbon dioxide + nitrogen dioxide + water
    C(s) + 4HNO3(l) CO2(g) + 4NO2(g) + 2H2O (l)

    Note: These acids oxidise to carbon to form carbon dioxide. It is the carbon dioxide gas that turns lime water milky.

    Self-evaluation Test 1.3

    1. Carbon reacts with carbon dioxide according to the equation below.
    CO2 (g) + C(s) 2CO (g)
    Identify the substance that loses oxygen and the one that gains oxygen from the reaction above.

    2. Carbon reacts with oxygen according to the equations below:
    i. C(s) + O2(g) CO2 (g) + heat
    ii. 2C(s) + O2(g) 2CO(g) + heat
    (a) What condition is required in (i) and (ii) above?
    (b) What does heat in the equations indicate about the two reactions?

    3. Write balanced chemical equations for the reactions that take place when:
    i. Carbon reacts with concentrated nitric acid.
    ii. Carbon reacts with concentrated sulphuric acid.

    1.4 Inorganic compounds of carbon

    Carbon can exist as a free element or in combined state. Commonly known inorganiccompounds of carbon include: oxides of carbon, carbonates and hydrogen carbonates.

    a) Oxides of carbon

    Research activity

    1. Using textbooks in library or from the internet, research on the physical properties of carbon oxides.
    2. Present your findings to the class.
    3. Compare your findings with the ones given below.

    The two oxides of carbon are carbon dioxide and carbon monoxide.

    i. Carbon monoxide

    This is a colorless and odorless gas. It is slightly denser than air. Carbon monoxideis produced by incomplete combustion (burning of carbon in limited oxygen supply) of various fuels. These include coal, wood, charcoal, fractions of oil and natural gas. It is a poisonous and very toxic gas. This is because when inhaled, it binds to haemoglobin blocking off oxygen. This can lead to death over prolonged exposure to the gas.

    ii. Carbon dioxide

    This is a colorless and odorless gas. Carbon dioxide is formed when fuels undergo complete combustion in sufficient supply of oxygen or air. It is a raw material required by plants for photosynthesis. Carbon dioxide is a greenhouse gas. It is produced in large quantities into the atmosphere through human activities in industries, motor vehicle emissions and burning of wastes. Excess carbon dioxide.

    in the atmosphere contributes to global warming that has resulted to climate change. The process of photosynthesis by plants helps to reduce the amounts of carbon dioxide in the atmosphere. This explains why people are advised to practise afforestration.



                                    My environment, my life!
    We should protect our existing forests and plant more trees. This will minimise effects of global warming.

    Self-evaluation Test 1.4

    1. Carbon monoxide is toxic. Explain.
    2. a) What is global warming?
    b) Discuss the causes of global warming.
    c) What steps should be taken to minimise the effects of global warming?

    b) Carbonates and hydrogen carbonates

    A carbonate is a salt formed when all the hydrogen atoms in carbonic acid are replaced by a metal or an ammonium ion. Carbonic acid is formed when carbon dioxide gas dissolves in water. Examples of metallic carbonates include:

    potassium carbonate, sodium carbonate, copper carbonate, zinc carbonate and magnesium carbonate.On the other hand, a hydrogen carbonate (also called bicarbonate) is formed when the hydrogen atoms in carbonic acids are partially replaced by a metal.The possible metallic hydrogen carbonates are sodium hydrogen carbonate and potassium hydrogen carbonate.

    1.5 Properties, preparation and testing for carbon dioxide, carbonates and hydrogen carbonates and their uses


    a) Laboratory preparation of carbon dioxide

    Activity 1.9

    Apparatus and reagents

    Flat-bottomed flask, dropping funnel, two wash bottles, delivery tubes, gas jars, cardboard covers, calcium carbonate (marble chips or limestone), dilute hydrochloric acid, distilled water, concentrated sulphuric acid.


    Procedure

    1. Arrange the apparatus as shown above.
    2. To the limestone or marble chips, add dilute hydrochloric acid dropwise.
    • Note down your observation.
    3. Collect the gas produced by downward delivery method (upward displacement of air) or over warm water.
    4. Collect a few jars of the gas and preserve them for the next experiment.

    Study questions

    1. Why is the gas collected by downward delivery method or over warm water?
    2. What is the role of water and concentrated sulphuric acid in the wash bottles?
    3. Suppose dilute sulphuric acid is used instead of hydrochloric acid in the above

    set up, what observation would you expect to make?
    4. Write a chemical equation for the reaction that takes place.

    Discussion corner!

    Discuss the study questions in pairs and write a report. Present your report to the rest of the class.

    I have discovered that...
    When dilute hydrochloric acid comes into contact with marble chips, effervescence occurs. This means a gas is produced.

    The facts

    Dilute hydrochloric acid reacts with marble chips (calcium carbonate) to form calcium chloride, water and carbon dioxide gas.
    Calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water
    CaCO3(s) + 2HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

    The gas is dried by passing it through concentrated sulphuric acid. It is then collected by downward delivery (upward displacement of air) as it is denser than air. Carbon dioxide can also be collected over warm water as it is only sparingly soluble in it. This is because the solubility of soluble gases in water decreases with increase in temperature.

    Water in the first wash bottle is used to remove sprays of hydrochloric acid (hydrogenchloride gas). The concentrated sulphuric acid in the set up is used to dry the gas. Anhydrous calcium chloride packed in a U–tube can also be used to dry carbon dioxide.

    All mineral acids react with metal carbonates to yield carbon dioxide gas. However,a reaction between dilute sulphuric acid and calcium carbonate is not feasible for preparation of carbon dioxide. This is because the reaction does not go to completion due to formation of an insoluble salt i.e. calcium sulphate that coats the rest of the carbonate preventing any further reaction.

    b) Test , properties and uses of carbon dioxide

    Activity 1.10

    Apparatus and reagents

    Nitric acid, gas-jars full of carbon dioxide, test-tubes, magnesium ribbon, a pair of tongs, candle, blue litmus paper, two water troughs, beehive shelves, calciumhydroxide solution, sodium hydroxide solution, dilute nitric acid, deflagrating spoon

    Procedure

    1. Examine the gas jar full of carbon dioxide.
    • Note down the colour of the gas and its smell.
    2. Pass carbon dioxide through a test-tube containing calcium hydroxide solution for a short time as shown below.


    Note down your observations.
    3. Continue passing more of the gas into the same test-tube until there is no further change.
    • Note down your observations.
    4. Pass carbon dioxide into a test-tube containing distilled water and then dip a blue litmus paper into the resulting solution.
    • Note down the colour change on the blue litmus paper.
    5. Invert a gas-jar full of carbon dioxide over a burning candle as shown below.
    • Note down what happens.
    6. Plunge a piece of burning magnesium ribbon quickly in a gas-jar full of carbon dioxide as shown below. • Note down your observations.

    7. Add about 50 cm3 of dilute nitric acid to the mixture in the gas jar.
    • Note down your observations. • Write a chemical equation for the reaction.
    8. Invert a gas jar full of carbon dioxide in a trough containing sodium hydroxide solution as shown below. • Shake the gas jar and wait for some time.

    9. Repeat procedure (8) above using water instead of sodium hydroxide.
    • Note any water level changes inside the gas jar.
    • What happens to the level of sodium hydroxide solution?
    • Write chemical equations for the reactions that occur.

    Study questions

    1. Write balanced chemical equations for the following:
    (a) Dissolving little amount of carbon dioxide into calcium hydroxide solution.
    (b) Dissolving excess carbon dioxide gas into calcium hydroxide solution.
    2. Explain why a solution of carbon dioxide in water turns blue litmus paper red.
    3. Write a balanced chemical equation for the reaction between carbon dioxide and water.
    4. Complete the chart below by finding the names of the compounds that correspond to A-D.


    5. Carbon dioxide does not support burning but magnesium continues to burn in its presence. Explain.

    Discussion corner!

    • Make groups of five and discuss the study questions above.
    • Write your suggested answers in your note books and submit to the teacher for evaluation.

    I have discovered that...
    Carbon dioxide is a colourless gas that does not smell. When a little amount of carbon dioxide is dissolved in calcium hydroxide solution, a white precipitate forms. However, the precipitate dissolves on addition of excess carbon dioxide. A solution of carbon dioxide in water does not completely turn blue litmus paper red. Carbon dioxide extinguishes a burning candle. However, magnesium continues to burn in it. When a gas jar containing carbon dioxide is inverted over a beaker containing sodium hydroxide solution, the level of sodium hydroxide in the gas jar rises.

    The facts

    The following are some of the properties of carbon dioxide.

    (a) Physical properties

    • Carbon dioxide is a colorless and odorless gas
    .• It is denser than air. That is why it can be collected by downward delivery.
    • It is slightly soluble in water.

    (b) Chemical properties

    (i) Reaction of carbon dioxide with calcium hydroxide

    A burning magnesium ribbon continues to burn inside a gas jar full of carbon dioxide. Black carbon particles form on the sides of the jar, together with white ash of magnesium oxide.

    The addition of dilute nitric acid enables the black carbon particles to be seen.

    more clearly because the acid reacts with magnesium oxide forming magnesium nitrate solution and water. Although carbon dioxide does not support combustion,magnesium continues to burn in it. The heat produced by the intense burning of magnesium is sufficient to decompose the carbon dioxide into constituent elements:carbon and oxygen. This oxygen produced supports the continued combustion to magnesium oxide.
    Carbon dioxide      heat         carbon + oxygen
                            
    CO2(g)     heat       C(s) + O2(g)
                
    Magnesium + oxygen magnesium oxide
    2Mg(s) + O2(g) 2MgO(s)

    The general equation for the reaction can then be written as follows:
    Magnesium + carbon dioxide magnesium oxide + carbon
    2Mg(s) + CO2(g)2MgO(s) + C(s)

    Magnesium being more reactive than carbon displaces it from oxygen.

    Uses of carbon dioxide

    Research activity

    1. Using textbooks and from the internet, research on the uses of carbon dioxide then write down your findings.
    2. Discuss your findings with a friend.
    3. Compare your findings with the ones outlined below.

    1. Carbon dioxide is used in the manufacture of sodium carbonate used in baking of cakes and bread among others. It is also an ingredient in some health salts to relieve constipation.

    2. Carbonated drinks

    Carbon dioxide is used as preservative in the production of mineral water and carbonated drinks like coca-cola (also called aerated or effervescence drinks).Soda-water is a solution of carbon dioxide in water under pressure. It is later sweetened, flavoured and sometimes coloured. The dissolved carbon dioxide gives it a pleasant taste. Figure 1.20 shows a girl opening a soda bottle. What do you think will happen to the contents of the bottle when the bottle top is removed?


    3. As a refrigerant

    Solid carbon dioxide commonly known as dry ice is a preferred refrigerant to ordinary ice (solid water) see figure 1.21.This is because it sublimes at room temperature forming gaseous carbon dioxide and therefore leaves no residue unlike ordinary ice.


    4. Fire extinguishers

    Some fire extinguishers contain sodium hydrogen carbonate or sodium carbonate solution and sulphuric acid. When mixed by inversion or pressing the plunger,carbon dioxide is produced due to the reaction between carbonate with sulphuric acid. In Fig 1.22 a person is trying to contain fire by using carbon dioxide fire extinguisher.


    Sodium hydrogen carbonate + sulphuric acid sodium sulphate + carbon dioxide + water
    NaHCO3 (aq) + H2SO4 (aq) Na2SO4 (aq) + 2CO2 (g) + 2H2O(l)

    Sodium carbonate + sulphuric acid sodium sulphate + carbon dioxide + water
    Na2CO3 (aq) + H2SO4 (aq) Na2SO4 (aq) + CO2 (g) + H2O (l)

    high pressure is set which forces a froth of carbon dioxide through the jet.
    It is directed onto the burning substances. In fire extinguishers that have carbon dioxide stored under pressure, the carbon dioxide is released through a nozzle.
    Carbon dioxide does not support combustion and since it is denser than air,it forms a blanket on the burning substance thereby cutting off the supply of oxygen from the air.

    5. Making rain during drought or in areas of little rain In making rain, dry ice (solid carbon dioxide) is spread in the clouds to accelerate the condensation process.Small aircrafts are sometimes used to spread dry ice in the sky.


    6. It is also used to transfer heat energy from certain types of nuclear reactors.
         Further activity
    Check the type of fire extinguishers you have in your school. Inquire from your teacher how they are used.

                             Self-evaluation Test 1.5


    1. State two physical properties of carbon dioxide gas.
    2. Write equation to show how sodium hydroxide reacts with carbon dioxide.
    3. Describe what happens when carbon dioxide is passed through a solution of calcium hydroxide (lime water). Use equations in your explanations.
    4. Carbon dioxide was passed through burning magnesium. State what was likely to be observed. Write equation for the reaction.
    5. The product in (4) was dissolved in water and a red litmus paper dipped into the solution formed.
    (a) State what was observed.
    (b) Write balanced equations for the reaction.

    1.6 Preparation, properties and test for carbonates and hydrogen carbonates

    Activity 1.11

    Apparatus and reagents

    Calcium chloride, lead (II) nitrate, sodium carbonate, filter funnel, filter paper, conical flask, beakers, stirring rod, spatula, measuring cylinder.

    Procedure

    1. Place 5 g of calcium chloride in a beaker and add to it 20 cm3 of water. Stir until all the solid dissolves.
    2. Put 5 g of sodium carbonate in another beaker and add to it 20 cm3 of distilled water, stir the mixture to form a uniform solution.
    3. Add the sodium carbonate solution to the calcium chloride solution and stir the mixture. Allow to settle. • Note down your observations.
    4. Filter off the precipitate obtained and wash it with distilled water.
    5. Dry the precipitate between filter papers.
    6. Repeat the procedure steps 1-5 using lead (II) nitrate in place of calcium chloride.



    Study questions

    1. What observation did you make when sodium carbonate solution was added to :
    i. Calcium chloride solution?
    ii. Lead (II) nitrate solution?
    2. Write balanced chemical equations for the reactions that occurred during the experiments.
    3. Which other solution can be used in place of sodium carbonate solution?

    Discussion corner!

    • In your study groups, discuss the results of these two experiments and the study questions above.
    • Write a joint report on your discussions and present it to the class.

    I have discovered that...

    When sodium carbonate solution is added to calcium chloride solution, a white precipitate forms. A similar precipitate is formed when lead (II)nitrate solution is mixed with sodium carbonate.

    The facts

    Sodium carbonate reacts with calcium chloride to form sodium chloride solution and insoluble calcium carbonate (white precipitate). Similarly, sodium carbonate reacts with lead (II) nitrate solution to form sodium nitrate solution and an insoluble lead (II) carbonate (white precipitate). The equations for the reactions are as follows:
    CaCl2(aq) + Na2CO3 (aq) CaCO3(s) + 2NaCl (aq)

    Pb(NO3)2(aq) + Na2CO3 (aq) PbCO3(s) + 2NaNO3 (aq)

    Filtering is done to separate the solid carbonate from the rest of the solution.The residue is then dried between filter papers or placed outside in the sun. Other carbonates such as magnesium carbonate and zinc carbonate can be prepared using the same method.

    b) Preparation of metal hydrogen carbonate

    Activity 1.12

    Apparatus and reagents

    Conical flask, carbon dioxide generator, delivery tubes, cold concentrated solution of sodium hydroxide, filter funnel, filter papers.

    Procedure

    1. Bubble carbon dioxide in excess through a cold concentrated solution of sodium hydroxide for 10 minutes.
    2. Filter the resulting mixture.
    3. Leave the residue undisturbed for sometime.
    4. Collect the crystals formed.

    Study questions

    1. What do you observe when carbon dioxide is bubbled through sodium hydroxide solution for 10 minutes?
    2. Write the equation for the reaction that occurs.

    Discussion corner!

    In groups of four discuss the study questions above and present your answers to the whole class.

    I have discovered that...
    When carbon dioxide is bubbled through sodium hydroxide for some time, a white precipitate is formed. After the residue is filtered and dried, crystals of a salt are formed.

    The facts

    In the laboratory, sodium hydrogen carbonate is prepared by passing carbon dioxide through cold concentrated solution of either sodium hydroxide or sodium carbonate. The equations for the reactions that take place are as follows:

    2NaOH (aq) + CO2(g) Na2CO3(aq) + H2O(l)

    Na2CO3 (aq) + H2O (l) + CO2 (g) 2NaHCO3(aq)

    The Solvay process

    Since the 1890s, sodium carbonate has been produced using Solvay process.Solvay process is a continuous process whereby limestone (CaCO3) is used to produce carbon dioxide, which then reacts with ammonia dissolved in brine (concentrated sodium chloride) to produce sodium carbonate.

    The steps in solvay process are:
    • Brine purification
    • Sodium hydrogen carbonate formation
    • Sodium carbonate formation
    • Ammonia recovery

    Environmental issues involved in production of sodium carbonate
    • Solid wastes such as calcium chloride, sand, clay and unburnt calcium carbonate.
    • Air pollution- ammonia released to the atmosphere is toxic.
    • Thermal pollution- due to release of large quantities of heat.Sodium carbonate has a number of uses but its most common use is in the production of glass.


    c) Chemical properties of carbonates and hydrogen carbonates

    i) Action of heat on carbonates and hydrogen carbonates

    Activity 1.13

    Apparatus and reagents

    Test-tubes, spatula, source of heat, calcium hydroxide solution, distilled water, red and blue litmus papers, copper (II) carbonate, magnesium carbonate, zinc carbonate, lead (II) carbonate, sodium carbonate, potassium carbonate, sodium hydrogen carbonate, calcium hydrogen carbonate, ammonium carbonate.

    Procedure

    1. Place a spatula full of the carbonate or hydrogen carbonate salts mentioned above in different test-tubes then heat.
    2. Test any gas formed with calcium hydroxide solution and moist litmus papers (red and blue).


    .Note down the colour of the residue when hot and cold for each residue.
    3. Record all your observations in a table like the one shown below.

    Table 1.1: Results for action of heat on carbonates and hydrogen carbonates

    Study questions

    1. Write balanced chemical equations for all the reactions that took place.
    2. Identify carbonates that behaved in a similar way.
    3. Which carbonate evolved a gas that turned red litmus paper blue?
    4. Which carbonate produced a residue that showed colour changes?
    5. Which carbonates were not affected by heat?

    Discussion corner!
    • In your groups, discuss the results of the experiment.
    • Classify the carbonates and hydrogen carbonates basing on similarity of their behaviour during and after heating.
    • Write a report on your conclusions and present it to the teacher for evaluation.

    I have discovered that...

    Sodium carbonate and potassium carbonate were not affected by heat. However, when the carbonates of magnesium, zinc, lead, copper as well as sodium hydrogen carbonate and calcium hydrogen carbonate were heated they produced a gas that formed a white precipitate with calcium hydroxide solution. The gas also turned moist blue litmus paper red. Ammonium carbonate produced two gases, one turned blue litmus paper red while the other turned red litmus paper blue.

    The facts

    Some carbonates decompose on heating while others do not. Carbonates of magnesium, zinc, lead and copper decompose on heating to produce respective metal oxide and carbon dioxide gas. Carbon dioxide gas produced forms a white precipitate with calcium hydroxide solution. The gas also turns moist blue litmus paper red as it is acidic. The metal oxides formed are of different colours as shown in the equations below.
                                         heat
    Copper (II) carbonatecopper(II) oxide + carbon dioxide
                         heat
    CuCO3(s)  CuO(s) + CO2(g)

    (green)                    (black)
     
    Magnesium carbonate     heat     magnesium oxide + carbon dioxide
                                        
    MgCO3(s)     heat     MgO(s) + CO2(g)
                    
    (white)                      (white)

    Zinc carbonate    heat     zinc oxide + carbon dioxide
                           
    ZnCO3(s)     heat     ZnO(s) + CO2(g)
                     
    (white)                         (white)
                                   heat 
    Lead (II) carbonatelead(II) oxide + carbon dioxide
                        heat
    PbCO3(s) PbO(s) + CO2(g)
    (white when hot)     (yellow on cooling)
                                   heat
    Calcium carbonate  calcium oxide + carbon dioxide
                        heat
    CaCO3(s) CaO(s) + CO2(g)
    (white)                       (white)

    Copper (II) oxide is black while oxides of magnesium and calcium are white. Zinc oxide is yellow when hot and white when cold. Lead (II) oxide is red-brown  when hot and yellow when cold. Carbonates of sodium and potassium are not affected by heat.

    Sodium hydrogen carbonate and calcium hydrogen carbonate decompose on heating to form a carbonate, water and carbon dioxide.
                                               heat
    Sodium hydrogen carbonate sodium carbonate + carbon dioxide + water
                                               heat
             2NaHCO3(s)                   Na2CO3(s)              +      CO2 (g) + H2O (l)
                                                    heat
    Calcium hydrogen carbonate         calcium carbonate + carbon dioxide + water
                                                          heat
                        Ca (HCO3)2(s        )       CaCO3(s)        +           CO2(g) + H2O(l)

    Note: When calcium hydrogen carbonate is strongly heated, calcium carbonate formed decomposes to form calcium oxide and carbon dioxide gas.Ammonium carbonate decomposes on heating to form ammonia gas, water and carbon dioxide.
                                        heat
    Ammonium carbonate  ammonia + carbon dioxide + water
                                        heat
               (NH4)2CO3(s)     2NH3(g) + CO2(g) + H2O(l)

    Ammonia and carbon dioxide gases are liberated at the same time. If moist red and blue litmus papers are put together at the mouth of the test-tube when ammonium carbonate is heated, ammonia gas is detected first (red litmus paper turns blue).

    ii) Action of dilute acids on carbonates and hydrogen carbonates

    Activity 1.14

    Apparatus and reagents

    Test tubes in a rack, spatula, calcium hydroxide solution, calcium carbonate,magnesium carbonate, zinc carbonate, lead (II) carbonate, copper carbonate,sodium carbonate, ammonium carbonate, sodium hydrogen carbonate, calcium hydrogen carbonate, dilute nitric acid, dilute sulphuric acid, dilute hydrochloric acid.

    Procedure

    1. In different test-tubes, put a spatulaful of the above carbonates and hydrogen carbonates and label them accordingly.
    2. Add dilute hydrochloric acid into each test tube and test the gas evolved with calcium hydroxide solution.


    .Note down the observation made in test tubes A and B.
    3. Repeat the experiment using sulphuric acid and then dilute nitric acid on all the carbonates and hydrogen carbonates
    .• Note down your observations in both test tubes A (reaction of acid + carbon/hydrogen carbonate and B reaction of gas produced + hydrogen carbonate solution).
    4. Record your observations in a table like the one shown below.

    Table 1.2 Reaction of carbonate or hydrogen carbonates with acid


    Study questions

    1. Explain the observations made when dilute hydrochloric acid and sulphuric acid were each added to lead (II) carbonate.
    2. What happened when dilute sulphuric acid was added to calcium carbonate solution?
    3. Write balanced chemical equations for all the reactions that took place.

    Discussion corner!
    • In your study groups, discuss the results obtained in this experiment
    .• Write your report on your findings and present it to the teacher for evaluation.

    I have discovered that...
    When an acid is added to any carbonate or a hydrogen carbonate, effervescence occurs, showing that a gas is produced. When the gas is dissolved in calcium hydroxide solution, a white precipitate is formed. However, on reacting lead (II) carbonate with either dilute sulphuric acid or hydrochloric acid, the reaction starts but stops after sometime. The same case happens when dilute sulphuric acid reacts with calcium carbonate.

    The facts

    Metal carbonates react with dilute mineral acids to produce a salt, water and carbon dioxide gas.

    Metal carbonate/hydrogen carbonate + dilute acid Salt + water + carbon dioxide
    The reaction between dilute sulphuric acid and calcium carbonate does not proceed to completion. This is due to formation of an insoluble coat (salt) of calcium sulphate that forms the impervious layer preventing any further reaction.Similarly, the reaction between lead (II) carbonate and dilute sulphuric acid as well as hydrochloric acid does not proceed to completion. This is due to formation of insoluble salts of lead (II) sulphate and lead (II) chloride respectively.

    iii) Solubility of carbonates and hydrogen carbonates in water

    Activity 1.15

    Apparatus and reagents

    Test tubes, distilled water, spatula, sodium carbonate, potassium carbonate,ammonium carbonate, sodium hydrogen carbonate, calcium hydrogen carbonate,calcium carbonate, zinc carbonate, lead (II) carbonate, magnesium carbonate.

    Procedure

    1. Place 0.5g of sodium carbonate in a test tube and add 5cm3 of distilled water.
    2. Label the test tube and shake its contents thoroughly.
    • Note down your observations.
    3. Repeat the experiment using the other carbonates and hydrogen carbonates provided and record your results in a table like the one shown below.

    Table 1.3: Solubility of carbonates and hydrogen carbonates in water


    Study questions

    1. Which carbonates are:
    (a) soluble in water?
    (b) insoluble in water?
    2. Which hydrogen carbonates are:
    (a) soluble in water?
    (b) insoluble in water?

    Discussion corner!

    • Discuss your results and the study questions above in your groups.
    • Present your conclusions to the whole class.

    I have discovered that...
    Carbonates of sodium, potassium and ammonium dissolve in water. All other carbonates do not dissolve in water. Sodium hydrogen carbonate also dissolves in water.

    The facts

    All carbonates are insoluble in water except ammonium carbonate, sodium carbonateand potassium carbonate. All hydrogen carbonates are soluble in water.

    Research activity

    Research and do a presentation on the uses of carbon compounds and their environmental issues under the following sub-headings.
    1. Carbon dioxide in fire extinguishers.
    2. Calcium carbonate in the manufacture of cement.
    3. Toxicity of carbon monoxide.
    4. Greenhouse effect and global warming.

                           Self-evaluation Test 1.6
    1. 2 cm3 of dilute nitric acid was added to 0.5g of zinc carbonate in a test tube.The gas produced was then dissolved in calcium hydroxide solution.
    (a) State the observations made.
    (b) Write stoichiometric equations for the reactions that occur.
    (c) Draw a set up of the apparatus that can be used to perform this experiment.
    2. The following set up was used to investigate the effect of heat on lead (II) carbonate.


    a)State the observations made in boiling tube A.
    b) What observation is made in test tube B?
    c) Write an equation for the reaction that occurs in test tube:
    i) A
    ii) B
    3. Name two carbonates that are:
    a) Soluble in water.
    b) Insoluble in water.
    4. A senior 3 student separately heated three solids A, B and C suspected to be carbonates. He tested the gases evolved (if any) with moist blue and red litmus papers. The following are the observations he made.


    a) Suggest the possible identities of carbonates A and C.
    (b) Explain the observations made with carbonate B.
    5. a) A solution reacts with crushed egg shells to give a gas that turns lime water milky. Name the gas that is evolved and give a brief explanation for this observation.
    b) Describe a chemical test that you could carry out to confirm the gas evolved in 5(a) above.
    c) Name the gas evolved when dilute hydrochloric acid reacts with sodium hydrogen carbonate. How is it recognised?
    6. Copy and complete the chemical equations below.
    a) H2SO4(aq) +Na2CO3(aq)
    b) HCl(aq) + Na2CO3(aq)
    c) HNO3(aq) + Na2CO3(aq)
    d) H2SO4(aq) + CaCO3(aq)
    e) HCl(aq) + NaHCO3(aq)
    7. Describe what would happen if the following carbonates were strongly heated:zinc carbonate, copper (II) carbonate and lead (II) carbonate. Write equations for the reactions that would take place.

    1.7 Environmental issues of carbon dioxide and carbon monoxide

    Both carbon dioxide and carbon monoxide gases have adverse effects on the environment.

    a) Environmental issues of carbon dioxide

    Activity 1.16

    Apparatus and reagents

    A large glass beaker or glass box, two small plastic beakers full of water, thermometer.

    Procedure

    1. Cover one of the small plastic beakers full of water with the big glass beaker or glass box as shown in Fig 1.27A below.
    2. Leave the other small plastic beaker with water in the open air as shown in Fig 1.27B.
    3. Put the two beakers under direct sunlight.


    4. At the end of the day, measure the temperature of the water in the two beakers.
    • Record the temperatures.

    Study question

    In which case did you record a higher temperature? Why?

    Discussion corner!

    In pairs, discuss the results of the experiment and present your findings to the rest of the class.

    I have discovered that...
    Water in the beaker that was covered with a glass box had a higher temperature at the end of the day.

    The facts

    Some gases are able to trap heat in the atmosphere keeping the earth’s surface warmer. Such gases are called ‘greenhouse’ gases. Carbon dioxide is an example of a greenhouse gas. Others include methane, chlorofluorocarbons and the oxides of nitrogen. As the amount of carbon dioxide in the atmosphere increases, the amount of heat retained by the atmosphere increases and the Earth becomes hotter. This is called greenhouse effect and it has led to the increase of temperatures in the world a phenomenon known as global warming.


    Human activities such as burning of fossil fuels and clearing forest land for agriculture and settlement have been blamed for the increase in amounts of carbon dioxide in the atmosphere. This has led to global warming which has in turn has caused unpredictable harsh climatic conditions of several regions. It is believed that

    with time, global warming will cause melting of polar ice. This will lead to coastal flooding that is dangerous to the lives of human beings, animals and plants.

    b) Environmental issues of carbon monoxide

    Carbon monoxide is formed through incomplete combustion of charcoal, coal, natural gas (methane) or firewood in poorly ventilated places. It also forms in improperly ventilated rooms in industries.Note: Carbon monoxide is not only a pollutant but also deadly, poisonous and toxic gas. Being colourless and odourless it is therefore not easy to detect. As mentioned earlier carbon monoxide competes with oxygen for the binding sites in haemoglobin (pigment that transports oxygen in the body).When carbon monoxide binds to haemoglobin molecules, little or no oxygen binds. This starves body tissues and cells of oxygen and may eventually cause death.

    Haemoglobin + carbon monoxide carboxyhaemoglobin

    The compound carboxyhaemoglobin is much more stable than oxyhaemoglobin (compound formed when oxygen binds with haemoglobin).This interferes with the transport of oxygen to various cells and tissues of the body.
    If noticed early, carbon monoxide poisoning can be reversed by giving the victim air enriched with oxygen and carbon dioxide. In severe cases of poisoning, a blood transfusion is essential.

                                   Self-evaluation Test 1.7
    1. Carbon monoxide is a poisonous gas. Explain.
    2. What is global warming?
    3. Name 3 greenhouse gases.
    4. Which of the following is not a major contributor to the greenhouse effect?
    (A) Carbon dioxide
    (B) Carbon monoxide
    (C) Chlorofluorocarbons
    (D) Methane gas
    (E) Nitrous oxide
    5. Which of the following is not likely a result of global warming?
    (A) Rising sea levels.
    (B) Increased agricultural productivity worldwide.
    (C) Increased storm frequency and intensity.
    (D) Increased coastal flooding.
    6. Which gas exists in the highest concentrations in the Earth’s atmosphere?

    1.8 Hard and soft water

    Activity 1.17

    Apparatus

    Distilled water, rain water, water from a borehole, bar soap, three basins, three pieces of dirty clothes.

    Procedure

    1. Put the various kinds of water from each of the three sources in different basins.
    2. Using soap, wash the dirty clothes seperately in each basin.
    • Note down the water that forms lather easily.

    Study questions

    1. In which water did soap lather easily?

    Discussion corner!

    Discuss your findings in pairs and come up with a conclusion.

    I have discovered that...
    Soap lathers easily in distilled water and rain water than in water from a borehole.

    The facts

    Hard water is water that contains salts of calcium and magnesium principally as bicarbonates, chlorides and sulphates. Hard water does not readily form lather with soap. Soft water on the other hand readily lather with soap. Distilled water and rain water are examples of soft water while water from borehole is an example of hard water.

    Water hardness is caused by dissolved calcium and magnesium salts. Hard water therefore contains calcium ions (Ca2+) and /or magnesium ions (Mg2+). Soft water on the other hand lacks these ions in solution.

    Temporary and permanent water hardness

    Research activity

    1. Using textbooks, journals from the library and from internet, research on the causes and how to remove temporary and permanent water hardness.
    2. Write a report on your findings and present it in class.
    3. Compare your findings with the ones outlined below.

    Water hardness can either be temporary or permanent.Temporary water hardness is caused by dissolved calcium hydrogen carbonate and magnesium hydrogen carbonate.Permanent water hardness on the other hand is caused by dissolved calcium sulphate and magnesium sulphate.

    Techniques of removing water hardness

    Activity 1.18

    Apparatus

    Water from a borehole or river, source of heat, boiling pot, calcium hydroxide,sodium carbonate, three basins and soap, pieces of dirty clothes.

    Procedure

    1. Divide the water from a borehole into three portions.
    2. Put one portion in a boiling pot then heat and cool, transfer to basin labelled A.
    3. Put the remaining portions into two basins, label them as B and C.
    4. To basin B add two spoonful of calcium hydroxide and mix thoroughly.
    5. To basin C add two spoonful of sodium carbonate and mix thoroughly.
    6. Add some soap in basin A, B and C. Try washing the dirty clothes in each portion. What happens?

    Study questions

    1. In which water did soap lather easily?
    2. Suggest the methods used above to remove water hardness.3. In which method was permanent water hardness removed?

    Discussion corner!

    Discuss your findings in pairs and come up with a conclusion.

    a) Removing temporary water hardness

    This type of hardness can be removed using the following methods:
    • Boiling
    • Distillation
    • Addition of calcium hydroxide
    i) Boiling
    As mentioned earlier, boiling removes the temporary water hardness by decomposing the dissolved calcium hydrogen carbonate or magnesium hydrogen carbonate to their respective carbonates i.e.

    Mg(HCO3)2 (aq)MgCO3(s) + CO2(g) + H2O(l)
    Ca(HCO3)2 (aq) CaCO3(s) + CO2(g) + H2O(l)
    ii) Distillation

    Distillation removes all dissolved solids. This method is used to prepare the distilled water we use in our chemistry laboratory. However, this method of water softening is too expensive to be used on a large scale.
    iii) Addition of calcium hydroxide

    When calcium hydroxide solution is added to temporary hard water, precipitation of the calcium and magnesium carbonate occurs. Calcium and magnesium ions are therefore removed from solution:

    Ca (HCO3)2(aq) + Ca(OH)2(aq) 2CaCO3(s) + 2H2O(l)Mg
    (HCO3)2(aq) + Ca(OH)2(aq) MgCO3(s) + CaCO3(s) + 2H2O(l)

    Note: It is important to add just enough of calcium hydroxide solution because when in excess, it would itself cause hardness.

    b) Removing permanent water hardness

    It is important to note that all methods used to remove permanent water hardness can as well remove temporary water hardness. Some of these methods include:

    i) Addition of washing soda (sodium carbonate)

    When sodium carbonate is added to hard water containing dissolved calcium sulphate and magnesium sulphate salts, their corresponding carbonates are precipitated.

    CaSO4(aq) + Na2CO3(aq) CaCO3(s) + Na2SO4(aq)
    MgSO4(aq) + Na2CO3(aq) MgCO3(s) + Na2SO4(aq)

    Notice that some sodium salts remain in solution but this does not affect the water in any way.

    ii) Use of ion exchange resins

    This involves passing the hard water through a column filled with a particular type of material (sodium permutit) which contains sodium ions. The calcium and magnesium ions are removed and remain in the column while the sodium ions in the column replace the calcium ions in the water.Note: Sodium ions in solution do not cause water hardness.

    Na2X (aq) + Ca2+ (aq) CaX(s) + 2Na+(aq)
    Na2X (aq) + Mg2+ (aq) MgX(s) + 2Na+(aq)


    When all the sodium ions on the permutit column have been exchanged with calcium or magnesium ions i.e. resin now contains ions of calcium and magnesium,it is said to be discharged. The permutit, can be “regenerated” or “recharged” by passing concentrated sodium chloride solution (brine) through it. Anotherion–exchange process takes place.


    Research activity

    1. Using textbooks and the internet, find out the advantages and disadvantages of hard water.
    2. Discuss your findings with a friend.
    3. Write a report and present it in class.
    4. Compare the advantages and disadvantages of hard water with those of softwater.

    Table 1.4 Advantages and disadvantages of hard water


                           Self-evaluation Test 1.8

    1. Differentiate between hard water and soft water.
    2. Write balanced chemical equations to show how water hardness is removed using the following methods:
    (a). Boiling
    (b). Use of washing soda
    3. Hard water is preferred for drinking while soft water is preferred for washing clothes. Explain.
    4. The type of water hardness that can be removed by boiling is called___________.
    5. Water which forms scum with soap is called ___________.

    1.9 The carbon cycle

                  Discussion corner!

    Study questions

    1. Why do plants take in carbon dioxide and release oxygen during the day.
    2. Why do animals take in oxygen and release carbon dioxide as a waste product.
    3. When we burn wastes, what is produced?
    4. What happens when you burn fossil fuels like coal, petroleum and diesel?

    • Discuss the above study questions in groups.
    • Summarise your findings in a report and present it to the rest of the class.Your report should include a cycle showing how carbon is recycled on Earth.

    The facts

    Some natural processes release carbon and its compounds into the atmosphere while other processes utilise this carbon and its compounds.Carbon dioxide is the compound through which carbon is recycled making it available to living things. The amount of carbon dioxide in the atmosphere remains fairly constant due to a delicate balance between processes that release the gas in to the atmosphere and those that remove the gas from the atmosphere. The circulation of carbon compounds in nature is called carbon cycle.

    a) Processes that add carbon dioxide to the atmosphere

    i. Combustion: This involves burning of petroleum, coal, wood, charcoal, wax or any other organic compound.

    ii. Respiration: During this process, sugar (glucose) obtained from food taken by animals is converted to release energy needed by the body while carbon dioxide and water are given out as by products. Carbon dioxide is exhaled to the atmosphere.

    iii. Decay of plants and animals: Any carbon (in plants and animals) is converted to carbon dioxide during the decaying process if sufficient oxygen supply is available.

    iv. Making of beer and wine: Yeast converts sugar to ethanol and carbon dioxide. This process is called fermentation. It is summarised using the equation below.

    sugar ethanol + carbon dioxide
    C6H12O6 (aq) 2C2H5OH (aq) + 2CO2 (g)

    The carbon dioxide formed in this process finally finds its way into the atmosphere.


    b) Processes that remove carbon dioxide from the atmosphere

    These include:
    i. Dissolution in water: Carbon dioxide dissolves in rivers and lakes to form weak carbonic acid. The carbonic acid reacts with calcium carbonate and magnesium carbonate to form calcium hydrogen carbonate and magnesium hydrogen carbonate respectively.

    carbon dioxide + watercarbonic acid
    CO2(g) + H2O (l) H2CO3(aq)
    Calcium carbonate + water + carbon dioxide calcium hydrogen carbonate
    CaCO3(s) + H2O (l) + CO2 (g) Ca (HCO3)2(aq)
    Magnesium carbonate + water + carbon dioxide magnesium hydrogen carbonate
    MgCO3(s) + H2O (l) + CO2 (g) Mg(HCO3)2(aq)

    ii.Photosynthesis: In this process, green plants use carbon dioxide to make sugar using sunlight as a source of energy.

    Carbon dioxide + water sugar + oxygen
    6CO2(g) + 6 H2O (l) C6H12O6(aq) + 6O2(g)

    iii. Reaction with calcium hydroxide: Natural calcium hydroxide slowly reacts with carbon dioxide from the air to form calcium carbonate.
    Calcium hydroxide + carbon dioxide calcium carbonate + water
    Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(l)

                            Self-evaluation Test 1.9
    1. Carbon is a necessary evil. Explain.
    2. If there is fire in the forest, there will be less___________ in the environment.(oxygen, carbon dioxide)
    3. Why is it important to maintain a balance of carbon dioxide in the atmosphere?
    4. Describe two processes that remove carbon dioxide from the atmosphere.
    5. Carbon cycle is the movement of carbon to the Earth by the process of ___________ and ______________.


    Unit summary 1

    • Allotropy is the existence of an element in more than one form in the same physical state.
    • Diamond and graphite are the two allotropes of carbon.
    • The uses of diamond and graphite are linked to the type of bond joining atoms and layers together and the presence or absence of delocalized electrons
    .• Carbon can exist as a free element or in compounds like oxides of carbon,carbonates and hydrogen carbonates.
    • Carbon acts as a reducing agent by removing oxygen from compounds like iron oxide, concentrated nitric acid and concentrated sulphuric acid
    .• In the laboratory, carbon dioxide is prepared by reacting a metal carbonate with any dilute mineral acid that forms a soluble salt
    .• The reaction between dilute sulphuric acid and calcium carbonate cannot be used to prepare carbon dioxide as the reaction does not proceed to completion.
    • Too much carbon dioxide in the atmosphere causes global warming
    .• Carbon monoxide is a very poisonous and toxic gas.
    • Water that contains dissolved calcium and magnesium ions is called hard water. Soft water lacks these ions
    .• Soft water lathers easily with soap. Hard water does not
    .• The whole process by which carbon dioxide circulates in the atmosphere is referred to as the carbon cycle.

    Test your Competence 1

    1. ______________ and ___________are two the crystalline allotropes of carbon while__________and ________are the non-crystalline forms of carbon.
    2. The flow-chart below shows some processes involving some compounds of carbon. Use it to answer the questions that follow.


    (a) Name the white precipitate A, solution B and gas C.
    (b) Describe a test that can be used to identify carbon dioxide in the laboratory.3. Akaliza set up an experiment as shown below.



    (a) What observation do you expect her to make in each test tube at the end of the experiment?
    (b) Write balanced equations that would lead to the observations expected in both test tubes.
    (c) Supposing she used sodium carbonate instead of copper carbonate, what observations would you expect her to make? Explain.
    (d) Write a balanced chemical equation for the decomposition of ammonium carbonate by heat.
    4. ___________ is a gas formed when carbon burns in enough oxygen supply while ___________ is formed when carbon burns in inadequate oxygen supply.
    5. Zinc oxide is ___________ in colour when hot and ___________ when cold while lead oxide is ___________ in colour when hot and ___________ when cold.
    6. Natural resources like fossil fuels are important and also harmful if not well utilised. Explain.
    7. The following are forms of amorphous carbon. Which one is not?
    A. Lamp black
    B. Charcoal
    C. Soot
    D. Graphite
    8. Hard water does not lather easily and hence wastes a lot of soap. Explain using chemical equations how you can convert hard water to soft water and use it in washing clothes.
    9. The diagram below represents part of the carbon cycle.


    Name processes A, B, C, D and E.
    10. Explain how you can minimise emission of greenhouse gases in your locality.
    11. Planting large numbers of trees results in ___________.
    A. A decrease in oxygen production.
    B. An increase in carbon dioxide production.
    C. A decrease in carbon dioxide built up in the atmosphere.
    D. An increase in oxygen use.

    12. Carbon dioxide is a gas produced by _______________ and _______________ processes.
    13. Identify and circle the raw materials required by plants during photosynthesis in the grid below.

    14. Identify and rewrite the correct word using the jumbled words given in the table below. Use the clues provided.