Topic outline

  • Forum: 1File: 1
  • 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.






    • Key unit competency

      After studying this unit, I should be able to relate the properties of nitrogen and its compounds to their uses, describe how some compounds of nitrogen are prepared and discuss the related environmental issues.

      Learning objectives

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

      • Describe the physical and chemical properties of nitrogen and its uses.

      • Describe the impact of nitrogen compounds on the environment.

      • Describe the industrial preparation of nitric acid and state its uses and those of some nitrates

      .• Prepare and collect ammonia gas and state its uses.

      • Interpret the effect of ammonia solution on different cations.

      • Develop observation skills during the experiments on addition of ammonia solution on different cations.

      • Develop self-confidence, a culture of working as a team and respect procedures in all experiments.

      • Appreciate the dangers posed by nitrogen compounds to the environment

      .• Carry out research on protection of the environment and hence protect natural resources.

      Mind teaser

      Like carbon, nitrogen is also important in our lives. Study these pictures carefully.


      How is nitrogen used in each picture above? What is the significance of that? Which other uses of nitrogen do you know?

      Introduction

      In senior one, you studied nitrogen as one of the components of air. What is the percentage composition of nitrogen in air? What compounds do you know that contain nitrogen?These compounds find application not only in Chemistry but also in other disciplines like Agriculture, Biology and Physics. For example,ammonia, nitric acid and some fertilisers are examples of compounds that contain nitrogen. It is therefore important to study how we can prepare nitrogen and use it for various purposes.


      2.1 Industrial isolation of nitrogen

      Activity 2.1

      1. Carry out a research using textbooks or from the internet about industrial isolation of nitrogen from other gaseous components. Use the sites http://www.fuseschool.org and http://www.youtube.com to do the research.

      Study questions

      1. Is air a mixture or a compound?

      2. Components of a mixture can be separated by physical means. Do you agree?

      3. What is the percentage composition of nitrogen in air?

      Discussion corner!

      Discuss the above study questions in your groups. Conclude whether it is possible to separate the components of air by physical means.

      The facts

      Air is a mixture of gases. The components of air include nitrogen, oxygen, water vapour, carbon dioxide, noble gases and dust particles. These components can be separated by physical means. As earlier mentioned, nitrogen forms about 78% of air by volume.

      Nitrogen can be isolated from air by fractional distillation of liquid air. This process involves eliminating other components of air hence remaining with nitrogen. It is a four step process as discussed below:

      (a) Purification of air

      This is the initial step. Air is purified to remove water vapour, dust and carbon dioxide. This is done by first passing air through filters to remove dust partic

      The dust-free air is then passed through driers to remove water vapour and finally it is passed through concentrated solution of sodium hydroxide. Sodium hydroxide absorbs carbon dioxide from the air. Note: It is important to remove water vapour and carbon dioxide before liquefaction to prevent the two from solidifying and blocking the pipes. Carbon dioxide and water vapour may also be removed by cooling the air to -80°C.

      (b)Compression of air

      This is the second stage whereby the remaining gas (mainly mixture of nitrogen, oxygen and noble gases) are compressed at very high pressures (200 atmospheres). This produces a lot of heat and so the compressed gases must immediately be cooled by passing them through cooling coils in a tank. The cooled compressed air is then allowed to expand by passing it through a narrow valve that then expands it rapidly. This makes its temperature to drop sharply. It is then returned to the compressor and the process repeated several times. After this repeated compression and expansion of air, cooling is done which liquefies air into a pale-blue liquid at -200°C. Neon and helium do not liquefy at this temperature and are hence removed at this stage.

      (c) Fractional distillation

      This is the third stage. The liquid air obtained is a mixture of colourless nitrogen (boiling point = -196°C), pale-blue oxygen (boiling point = -183°C) and argon (boiling point = -186°C). Air is thus passed through a fractionating column. The temperature of liquid air in the fractionating column is slowly raised. Nitrogen boils at -196°C and so distills first and is collected at the top of the fractionating column. Argon with a boiling point of -186°C and oxygen with a boiling point of -183 °C distills later and are collected from the lower part of the column in that order.


      (d) Storage of nitrogen

      The nitrogen obtained is liquefied and stored under pressure.

                              Self-evaluation Test 2.1

      1. Why is it possible to isolate nitrogen from air?

      2. What is the importance of the following during isolation of nitrogen from air:

      (a) Passing air through filters and concentrated sodium hydroxide solution?

      (b) Compression and expansion of air?

      (c) Fractional distillation of liquid air?

      3. In your study groups, discuss the industrial isolation of nitrogen gas from air and make a presentation to the rest of class.

      4. Air contains mainly _____________ gas with about ___________%. The main components of air are mainly separated by _______________process.

      5. Liquid air has three components X, Y and Z whose boiling points are -186°C, -183°C and -196°C respectively. When the liquid air is fed into a tall fractional distillation column and near its bottom warmed up slowly:

      (a) Which component will be collected from near the bottom of the fractional distillation column? Why? (b) Which component will be collected from the top of the fractional distillation column? Why?

      (c) Which component will be collected from the middle part of the fractional distillation column? Why?

      (d) What could the components X, Y and Z be in real situation?

      2.2 Laboratory preparation of nitrogen

      Activity 2.2

      Apparatus and reagents

      Round-bottomed flask or boiling tube, Bunsen burner/any source of heat, trough, gas jars, sodium nitrite, ammonium chloride.

      Procedure

      1. Make a solution of a mixture of sodium nitrite and ammonium chloride in water.

      2. Put the solution in a round-bottomed flask or boiling tube.

      3. Arrange the apparatus as shown in Fig. 2.4 below.


      4. Warm the flask slightly and then remove the heat.

      5. Record your observations in the flask and gas jar.

      Study questions

      1. Why was warming done only for a short time?

      2. Write balanced chemical equations for the reactions that take place.

      Discussion corner!

      Discuss your observations in groups and present your findings is class.

      I have discovered that...

      When the flask is warmed, bubbles are produced and this continues even after warming has been stopped. A colourless gas collects in the upper part of the gas jar.

      The facts

      Sodium nitrite reacts with ammonium chloride to form ammonium nitrite and sodium chloride.

      Sodium nitrite + ammonium chloride sodium chloride + ammonium nitrite

      NaNO2 (aq) + NH4Cl (aq) NaCl (aq) + NH4NO2 (aq)

      The ammonium nitrite is then decomposed to produce nitrogen gas.

      Ammonium nitrite nitrogen + water

      NH4NO2 (aq) N2(g) + 2H2O(l)

      The overall equation for the reaction:

      Sodium nitrite + ammonium chloride sodium chloride + water + nitrogen

      NaNO2 (aq) + NH4Cl (aq) NaCl(aq) + 2H2O(l) + N2(g)

      Note: If required dry, nitrogen gas can be passed through concentrated sulphuric acid. The gas is then collected in a syringe or by upward delivery.Nitrogen can also be prepared by direct heating of ammonium nitrite. However,when ammonium nitrite is in small quantities, the reaction can be explosive.

      a) Properties and uses of nitrogen

      Activity 2.3

      Apparatus and reagents

      Wooden splint, blue and red litmus papers, aqueous calcium hydroxide, magnesium ribbon, sulphur powder, six gas jars full of dry nitrogen, a pair of tongs, deflagrating spoon.

      Procedure

      1. Copy Table 2.1 in your notebook.

      2. Carry out the procedures 1 -7 and use your observations and conclusions to fill the table.

      Table 2.1: Investigating the properties of nitrogen


      Study questions

      Write chemical equations for the reactions that take place in procedure number 6.

                                     Discussion corner!

      In your study groups, discuss the observations made in each of the investigations.

      • Present your results to the rest of the class.

      I have discovered that...

      Nitrogen shows no observable reactions for most of the tests carried out.

      The facts

      The following are some properties of nitrogen gas.

      a) Physical properties of nitrogen

      • Nitrogen is colourless and does not smell (odourless).

      • It is almost insoluble in water and hence can be collected over water.

      • It is lighter than air and therefore is collected by upward delivery method.

      b)Chemical properties of nitrogen

      (i). Combustion

      Nitrogen neither burns nor supports combustion. That is why a glowing or burning splint goes off when introduced in nitrogen.

      (ii). Effect on litmus papers

      Nitrogen is neutral. Litmus papers therefore show no colour change in nitrogengas.

      (iii). Effect on aqueous calcium hydroxide

      Nitrogen has no effect on calcium hydroxide.

      (iv). Reaction with magnesium and sulphur

      Nitrogen reacts with burning magnesium to form magnesium nitride.

      Magnesium + nitrogen magnesium nitride

      3Mg(s) + N2(g) Mg3N2(s)

      Magnesium nitride reacts with water to form magnesium hydroxide solution and ammonia gas.

      Magnesium nitride + water magnesium hydroxide + ammonia
      Mg3N2(s) + 6H2O (l) 3Mg(OH)2(aq) + 2NH3(g)

      However, when burning sulphur is introduced into a gas jar of nitrogen, no reaction occurs. This is because the heat produced from burning sulphur is not sufficient to break the triple covalent bonds        (  NN) of the nitrogen molecules for the reaction to occur.

      Nitrogen is inert at room temperature. It does not react with other elements at room temperature. However, the high temperatures produced by burning metals like magnesium break down the triple covalent bonds of the nitrogen molecules to form single atoms of nitrogen which is now able to react with magnesium.
      NN 2N (very high temperature needed).

      Uses of nitrogen

      1. It is used in the Haber process to manufacture ammonia gas.
      2. It is used to freeze substances (coolant)due to its low boiling point of -196°C. It can be used to store body tissues that are required to last for a long period of time. For instance:
      • Tissues used in hospitals.
      • Bulls’ semen for artificial insemination.
      • It can also be used to mend leaking pipes. When liquid nitrogen is poured onto the pipe, it freezes the liquid inside while repair is done.
      3. It is used in food processing. When food is being packed, nitrogen is used to keep off oxygen and the food stays fresh for a long time because no oxidation can take place. Most bacteria cannot survive in oxygen.
      4. Because of its inert nature at low temperature, it is pumped into the ships’tanks that transport crude oil to remove traces of oxygen. This prevents dangerous explosions that could occur from crude oil vapours.

                                   Self-evaluation Test 2.2

      1. Nitrogen gas does not support combustion but a magnesium ribbon continues to burn in the presence of nitrogen. Discuss.
      2. Liquid nitrogen boils at -196°C. What is the importance of this?
      3. Give three physical properties of nitrogen gas that differentiate it from other gases.
      4. In order to attain electronic configuration of noble gas, nitrogen needs to ––––––.
      A. Lose one electron
      B. Gain two electrons
      C. Loose three electrons
      D. Gain three electrons

      2.3 Preparation, properties and uses of nitrogen compounds

      Compounds to be studied under this section include:
      • Nitrogen dioxide                • Ammonia                      • Nitric acid

      a) Nitrogen dioxide

      Activity 2.4

      Apparatus and reagents

      Thistle funnel or dropping funnel, flat-bottomed flask , delivery tube, 3 gas jars , cardboard cover , copper turnings, concentrated nitric acid.

      Caution: Nitrogen dioxide is poisonous. This experiment should therefore be performed in a fume cupboard or in an open space.

      Procedure

      1. Put copper turnings into a flat-bottomed flask.
      2. Arrange the apparatus as shown in Fig. 2.5.

      3. Add concentrated nitric acid to the copper turnings in the flask.
      4. Make your observations in the flask and the gas jar and record in your notebook.

      Study questions

      1. What observations did you make in the flask and the gas jar?
      2. Write a balanced chemical equation for the reaction that takes place.

                                               Discussion corner!
      • In your study groups, discuss the results of the experiment and the study questions above
      .• Share your findings with the rest of the class.

      I have discovered that...
      When concentrated nitric acid is added to copper turnings, effervescence occurs immediately producing red-brown fumes. A green solution is also formed in the flask.

      The facts
      Copper reacts with concentrated nitric acid to form a green solution of copper (II) nitrate, water and a red-brown gas (nitrogen dioxide)

      Copper + nitric acid copper (II) nitrate + water + nitrogen dioxide
      Cu(s) + 4HNO3 (l) Cu(NO3)2(aq) + 2H2O (l) + 2NO2 (g)

      More pure nitrogen dioxide can be prepared by heating a nitrate of a moderately reactive heavy metal. Lead (II) nitrate is preferred because unlike most nitrates,they do not contain water of crystallisation which would otherwise interfere with the preparation process since nitrogen dioxide is highly soluble in water.
      2Pb(NO3)2(s) 2PbO(s) + 4NO2 (g) + O2 (g)


      In this method, the lead (II) nitrate decomposes on heating to give a mixture of nitrogen dioxide and oxygen. The nitrogen dioxide is cooled by freezing in a mixture ice and salt forming a pale-yellow liquid called dinitrogen tetroxide (N2O4). Oxygen passes through as a gas and is collected over water. This way, the two gases are separated. Dinitrogen tetraoxide easily dissociates into nitrogen dioxide when the temperature is raised as per the equation below.

                                                           N2O4(l)    2NO2(g)
                                                   (Pale - yellow)                             (Red - brown)

      Properties of nitrogen dioxide

      a) Physical properties of nitrogen dioxide

      • It is a red-brown gas
      • It has an irritating pungent smell.
      • It is soluble in water forming an acidic solution
      2NO2 (g) + H2O(l) HNO3(aq) + HNO2(aq)
      • It is one and half times denser than air.
      • It is easily liquefied on cooling (boiling point is 21°C).

      b) Chemical properties of nitrogen dioxide

      Activity 2.5

      Apparatus and reagents

      Litmus papers (blue and red), magnesium ribbon, phosphorus, sulphur, copper turnings, a pair of tongs, wooden splints.

      Procedure

      1. Copy Table 2.2 in your notebook.
      2. Prepare nitrogen gas as shown in Activity 2.4 then carry out the following investigations 1-7. Record
      your observations and conclusions in the blank spaces provided.

      Table 2.2: Investigating the chemical properties of nitrogen dioxide



      Study questions
      1. Account for the observations in procedures 4 and 5.
      2. Explain the observations made in procedure 6.

                                        Discussion corner!

      • In your study groups, discuss the observations made in each procedure and the study questions above.
      • Draw appropriate conclusions and present your findings to the rest of the class.

      I have discovered that...
      Both glowing and burning splints goes off on exposure to nitrogen dioxide gas. Burning magnesium, phosphorus and sulphur continue to burn in the presence of nitrogen dioxide.

      The facts

      The following are the chemical properties of nitrogen dioxide.

      i. Combustion

      Nitrogen dioxide neither burns nor supports combustion. That is why it puts off both the glowing and the burning splints.

      ii. Reaction with water

      Nitrogen dioxide reacts with water to form nitric acid and nitrous acid.

      Nitrogen dioxide + water nitric acid + nitrous acid
      2NO2 (g) + H2O(l) HNO3(aq) + HNO2(aq)

      The gas is therefore acidic and so the solution formed turns wet blue litmus paper red.

      iii. Effects on burning elements

      Burning magnesium continues to burn in nitrogen dioxide forming magnesium oxide and nitrogen gas.

      Magnesium + nitrogen dioxide magnesium oxide + nitrogen gas
      4Mg(s) + 2NO2(g) 4MgO(s) + N2(g)

      The heat from burning magnesium decomposes this gas into its constituent elements nitrogen and oxygen.
      The oxygen produced supports the continued burning of magnesium.The same type of reaction takes place if vigorously burning phosphorus and sulphur are introduced into gas jar containing nitrogen dioxide.

      Phosphorus + nitrogen dioxide phosphorus pentoxide + nitrogen gas
      8P(s) + 10NO2 (g) 4P2O5(s) + 5N2 (g)

      Sulphur + nitrogen dioxide sulphur dioxide + nitrogen gas
      2S(s) + 2NO2(g) 2SO2(g) + N2(g)

      Nitrogen dioxide oxidises red-hot copper to copper (II) oxide while itself is reduced to nitrogen gas.

      Copper + nitrogen dioxide copper (II) oxide + nitrogen gas
      4Cu(s) + 2NO2 (g) 4CuO(s) + N2 (g)

      Note: Heated nitrogen dioxide easily reacts with burning substances because the heat decomposes it to nitrogen monoxide and oxygen.

      Nitrogen dioxide nitrogen monoxide + oxygen
      2NO2 (g) 2NO(g) + O2(g)

      Nitrogen monoxide is further decomposed to nitrogen and oxygen.

      Nitrogen monoxide nitrogen + oxygen
      2NO (g) N2(g) + O2(g)

      Finally the oxygen combines with element in given reaction to form their respective oxides.

      iv. Reaction with sodium hydroxide

      solutionNitrogen dioxide reacts with sodium hydroxide solution to form two salts that is sodium nitrate (NaNO3) and sodium nitrite (NaNO2). The two salts are formed because nitrogen dioxide when dissolved in water forms two acids; nitric acid (HNO3) and nitrous acid (HNO2).

      Nitrogen dioxide + water nitric acid + nitrous acid
      2NO2 (g) + H2O (l) HNO3 (aq) + HNO2 (aq)

      These acids though present in the same solution, separately react with sodium hydroxide solution to form the two salts.
      Sodium hydroxide + nitric acid sodium nitrate + water
      NaOH (aq) + HNO3 (aq) NaNO3 (aq) + H2O (l)

      Sodium hydroxide + nitrous acid sodium nitrite + water
      NaOH (aq) + HNO2 (aq) NaNO2 (aq) + H2O (l)

      The overall equation for the reaction between nitrogen dioxide and sodium hydroxide is therefore as shown below.
      Sodium hydroxide + nitrogen dioxide sodium nitrate + sodium nitrite + water
      2NaOH (aq) + 2NO2 (g) NaNO3 (aq) + NaNO2 (aq) + H2O(l)

      Uses of nitrogen dioxide

      1. It is used as an intermediary in the manufacture of nitric acid by Ostwald Process.
      2. It is an intermediary in the formation of very dilute nitric acid during lightning. This acid is washed into the soil, thereby increasing the amounts of nitrates in the soil.

                                        Health check!
      Nitrogen dioxide and the environment
      The exposure to high levels of nitrogen dioxide (NO2 ) may result in respiratory disorders. Nitrogen dioxide in flames the lining of the lungs which could lead to reduced lung immunity. This can cause further problems such as wheezing, coughing, colds, flu and bronchitis.
      Increased levels of nitrogen dioxide can have significant health risk on persons with asthmatic conditions because it can lead to more frequent and more intense attacks. Children with asthma and older people with heart diseases are most at risk.Visit: www.greenfacts.org for more information

                                         Self-evaluation Test 2.3

      1. Nitrogen dioxide dissolves in water to form two acids. Write the chemical equation for the reaction that takes place.
      2. Nitrogen dioxide neither burns nor supports combustion. However, burning magnesium continues to burn in the presence of the gas. Explain.
      3. Give one use of nitrogen dioxide.

      b) Ammonia

      Laboratory preparation of ammonia

      Activity 2.6

      Apparatus and reagents

      Round-bottomed flask, glass tubing, spatula, drying tower, several gas jars, mortar and pestle, cardboard cover, calcium hydroxide, ammonium chloride, calcium oxide, litmus papers (blue and red).

      Procedure

      1. Place one spatulaful of calcium hydroxide and one spatulaful of ammonium chloride in a mortar.
      2. Grind the mixture well using the pestle then put it in a round-bottomed flask.
      3. Arrange the apparatus as shown in fig 2.7.


      Caution: The flask should be set in a sloping position as shown in the diagram before the experiment starts.
      4. Heat the mixture.
      • Note down what you observe on the cooler parts of the flask and the delivery tube.
      5. Check whether the gas jar is full, by putting a moist red litmus paper at the mouth of the jar. It should turn blue if the gas jar is full.
      6. Collect five jars of this gas for the next activity 2.7.

      Study questions

      1. Why should the flask be in a sloping position?
      2. Identify the chemical labelled substance A in the Fig 2.7 and state its function.3. Name the method used to collect ammonia gas and suggest why it is preferred.

                                      Discussion corner!
      • In groups, discuss the above method of preparing ammonia gas and the study questions above.
      • Write a group report on the experiment and present it to the teacher for evaluation.

      I have discovered that...
      When the mixture is heated, water droplets are seen on the cooler parts of the flask and delivery tube. When a damp red litmus paper is placed at the mouth of the gas jar, it turns blue.

      The facts

      Calcium hydroxide reacts with ammonium chloride to form calcium chloride,water and ammonia gas. The equation for the reaction is as follows:

      Calcium hydroxide + ammonium chloride calcium chloride + water + ammonia
      Ca(OH)2(s) + 2NH4Cl(s) CaCl2(s) + 2H2O(l) + 2NH3(g)

      Substance A is calcium oxide. It is used to dry ammonia gas. Other drying agents like concentrated sulphuric acid and anhydrous calcium chloride would react with ammonia and hence cannot be used to dry the gas.

      Ammonia + sulphuric acid ammonium sulphate
      2NH3 (g) + H2SO4(l) (NH4)2SO4(aq)

      Ammonia + calcium chloride calcium ammonium chloride
      4NH3 (g) + CaCl2 (aq) CaCl2.4NH3(s)
                                                    (Solid complex)

      The flask should be in a slanting position to prevent drops of water vapour that condenses on the cooler parts of the flask from flowing back into the flask which would make the flask to crack.Ammonia is less dense than air and is hence collected by upward delivery (downward displacement of air).

      Properties of ammonia gas

      Activity 2.7

      Apparatus and reagents

      Four gas jars of ammonia, wooden splints, trough, glass rod, concentrated hydrochloric acid.

      Procedure

      1. Copy table 2.3 in your notebook.
      2. Carry out procedures 1-5 and use your results to fill the table.

      Table 2.3: Investigating the properties of ammonia gas



      Study questions
      1. Describe the smell of ammonia gas.
      2. Is ammonia an acidic or alkaline gas?


                                    Discussion corner!

      • In groups, discuss the study questions and properties of ammonia based on your observations in this activity.
      • Write a report and present it to the teacher for evaluation.

      I have discovered that...
      Ammonia is a colourless gas with a pungent choking smell. It reacts with hydrogen chloride forming white fumes. It is very soluble in water (1cm3 of water dissolves about 800 cm3of the gas).

      The facts

      Physical properties of ammonia

      1. Ammonia is a colourless gas.
      2. It has a pungent choking smell.
      3. It is very soluble in water. Water rises rapidly in a round-bottomed flask full of ammonia gas and almost fills it immediately if inverted in a trough containing water. This can be demonstrated in a fountain experiment as shown in the set up that follows.

      In the set up above, when water enters the flask, it dissolves so much ammonia that there is a partial vacuum in the flask. Water is then rapidly forced up the tube and enters the flask as a fountain.

      Further activity

      Carry out an experiment to demonstrate the fountain experiment. Arrange your apparatus as shown in the figure above.

      Chemical properties of ammonia gas

      1. Combustion
      Ammonia does not burn in air neither does it support combustion.

      2. Effect on litmus paper
      Ammonia turns moist red litmus paper blue. It is the common known alkaline gas.

      3. Reaction with hydrogen chloride
      It reacts with hydrogen chloride gas to form dense white fumes of ammonium chloride.

      NH3 (g) + HCl (g) NH4Cl(s)
                                           (dense white fumes)

      Note: This is the confirmatory test for ammonia.

      4. Reaction of ammonia with oxygen
      Ammonia reacts differently with air or oxygen. The reaction is even more different when a catalyst is introduced.

      Apparatus and reagents

      Open glass tube, delivery tube, glass beads or cotton wool, sources of oxygen and ammonia, wooden splint, cork/rubber bung.

      Procedure

      1. Arrange the apparatus as shown in the figure below.



      2. Ignite ammonia gas at the jet delivering ammonia before allowing oxygen in as shown above.
      3. Now adjust the rate of flow of both gases until a flame can be lit at the jet of ammonia tube.
      4. Record your observations in your notebook.

      Study questions

      1. What is the use of the glass beads or cotton wool in this experiment?
      2. What do you conclude from this experiment?

                                                  Discussion corner!

      Discuss the experiment results in your groups and answer the study questions above. Present your findings to the rest of the class.

      I have discovered that...
      Ammonia burns in presence of oxygen with a yellow-brown flame.

      The facts

      Ammonia burns in oxygen to form nitrogen gas and water.
      Ammonia + oxygen nitrogen + water
      4NH3 (g) + 3O2 (g) 2N2 (g) + 6H2O (l)

      The loosely packed glass beads or cotton wool helps to spread out the oxygen.

      Note: Ammonia gas neither burns nor supports burning. It can only burn in oxygen or air enriched with oxygen.

      Catalytic oxidation of ammonia

      Activity 2.9

      Apparatus and reagents

      Platinum coil, delivery tube, beaker, concentrated ammonia solution, source of oxygen, litmus papers (blue and red).
      Caution: The reaction can sometimes be explosive.

      Procedure

      1. Bubble oxygen through the concentrated ammonia solution for a few minutes.
      2. Lower a red-hot platinum coil into the mixture of oxygen and ammonia solution.
      Note: Concentrated ammonia solution produces ammonia gas.
      3. Record your observations in your notebook.
      4. Test the fumes produced using moist blue and red litmus papers.
      5. Dip blue and red litmus papers in the solution after the reaction.

      Study questions

      1. What is the color of the fumes that are seen coming out of the solution?
      2. What is the effect of these fumes on moist blue litmus paper?
      3. What is the effect of the resulting solution on blue litmus paper?
      4. What conclusions can you make from your observations?

                               Discussion corner!
      • Discuss the results of the experiment in your groups and answer the study questions above
      .• Present your discussion report to the rest of the class.

      I have discovered that...

      Red-brown fumes are produced when the reaction takes place. The fumes turn blue litmus paper red but have no effect on red litmus paper.

      The facts

      Ammonia reacts with oxygen in the presence of platinum (Pt) - Rhodium (Rh) catalyst to form nitrogen monoxide and water as shown in the equation below.

      Ammonia + oxygen nitrogen monoxide + water
      4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

      The nitrogen monoxide is further oxidised to form nitrogen dioxide gas characterised by red-brown fumes. This is an acidic gas.

      Nitrogen monoxide + oxygen nitrogen dioxide
      2NO (g) + O2 (g) 2NO2 (g)
                                            (Red-brown fumes)

      When nitrogen dioxide is dissolved in water, nitrous acid and nitric acid are formed.

      Nitrogen (IV) oxide + water nitric acid + nitrous acid
      2NO2 (g) + H2O (l) HNO3 (aq) + HNO2 (aq)

      Nitrous acid is further oxidised to form nitric acid.Nitric acid is the one that turns blue litmus paper red.

      Nitrous acid + oxygen nitric acid
      2HNO2 (aq) + O2 (g) 2HNO3 (aq)

      It is therefore possible to convert ammonia to nitric acid using platinum as a catalyst.

      Manufacture of ammonia by Haber process

      Research activity

      1. In groups and using textbooks, research on the industrial manufacture of ammonia by the Haber process.
      2. Compile a report then present it to the rest of the class.

      The facts

      The raw materials required for the manufacture of ammonia by the Haber process
      are:
      i.Nitrogen- obtained from air by fractional distillation of liquid air.
      ii.Hydrogen -obtained from natural gas (methane) or by electrolysis of acidified water or brine where electricity supply is cheap and readily available.

      Optimum conditions needed during manufacture of ammonia are:
      • A mixture of nitrogen and hydrogen in the ratio of 1:3 by volume.
      • A pressure of between 200– 250 atmospheres.
      • A temperature of between 400°C – 500C.
      • Iron catalyst in finely divided form and impregnated with aluminium oxide as
      a promoter of the catalyst.

      Reactions that take place

      When nitrogen and hydrogen react under the conditions mentioned above, only about 10% of ammonia is formed. The reaction is exothermic.

      Nitrogen + hydrogen     ammonia + heat
      N2(g) + 3H2(g)   2NH3(g) + heat



      The heat produced   during the reaction  is taken back to the heat exchanger together with the unreacted nitrogen, hydrogen  and the 10% ammonia formed. This helps to maintain the temperature needed for the reaction at 500°C. The ammonia produced and unreacted hydrogen and nitrogen are then taken to the liquifier or condenser where ammonia is liquified or dissolved in water. The unreacted nitrogen and hydrogen are re-circulated to the compressor.

      Aqueous ammonia

      Preparation of aqueous ammonia (ammonium hydroxide)

      Activity 2.10

      Apparatus and reagents

      Beaker, delivery tube, filter funnel, source of ammonia.

      Procedure

      1. From the preparation flask, pass ammonia gas into water via an inverted filter funnel as shown in the figure 2.13 for some time. The filter funnel should just dip into the water surface.


      2. Smell the aqueous solution in the beaker.
      3. Record your observations in your notebook.

      Study questions

      1. Why is the funnel preferred to a delivery tube in dissolving ammonia in water?
      2. Why is the ammonia gas used in this experiment not dried?

                          Discussion corner!
      • Discuss the procedure and observations made in groups.
      • Compile a report and present it to other class members.

      I have discovered that...

      After sometime, the water takes the smell of ammonia.

      The facts

      Ammonia gas dissolves in water to form aqueous ammonia or ammonia solution.Ammonia solution is also called ammonium hydroxide. The gas need not to be dried because water has been added to ammonia.

      Ammonia + water ammonium hydroxide
      NH3 (g) + H2O (l) NH4OH (aq)

      The funnel creates a large surface area for dissolution of ammonia as opposed to a delivery tube. In addition, if a delivery tube is used in place of the funnel, water would “suck back” into the hot preparation flask thereby cracking it.Dissolution of ammonia in water is a reversible reaction. This is why aqueous ammonia produces ammonia gas at any temperature. It decomposes on warming to produce a strong smelling gas which is ammonia.

      Properties of aqueous ammonia

      Reaction of ammonia with dilute acids

      Activity 2.11

      Apparatus and reagents

      Test tubes and test tube rack, 2M nitric acid, 1M sulphuric acid, 2M hydrochloric acid, methyl orange indicator, aqueous ammonia, labels.

      Procedure

      1. Place 2cm3 of 1M sulphuric acid, 2M hydrochloric acid and 2M nitric acid in three separate test tubes and label them accordingly.
      2. Add three drops of methyl orange indicator in each test tube.
      3. Add aqueous ammonia drop wise into each test tube until a permanent pink orange colour is formed.

      4. Record your observations.

      Study questions

      1. What name is given to the types of reactions that occur between aqueous ammonia and dilute acids?

      2. Write the equations for the following reactions between aqueous ammonia and dilute:
      a. sulphuric acid
      b. hydrochloric acid
      c. nitric acid

                         Discussion corner!
      • In your study groups discuss the observations made.
      • Compare your findings with those of other groups in the class.

                     I have discovered that...
      When ammonia solution is added to dilute acids, the solution turns yellow.

      The facts

      Aqueous ammonia neutralises dilute acids forming the corresponding ammonium salt and water only. The equations for the reactions that occur are as follows:

      H2SO4 (aq) + 2NH4OH (aq) (NH4)2SO4(aq)+ 2H2O(l)
      HCl (aq) + NH4OH (aq) NH4Cl(aq) + H2O(l)
      HNO3 (aq) + NH4OH (aq) NH4NO3(aq) + H2O(l)

      All ammonium salts are soluble in water and are normally prepared by neutralisation reaction. It is then followed by crystallisation of the salts.

      Uses of ammonia

      1. Large quantities of ammonia are used to make fertilisers.
      2. Ammonia gas is used in the manufacture of nitric acid.
      3. Liquid ammonia is used as a refrigerant in large-scale refrigerating plants and factories.
      4. Ammonia solution is used as a solvent during cleaning (in laundries).
      5. Ammonia is used in the manufacture of ammonium salts such as ammonium

      chloride, which is used in dry cells and ammonium carbonate used in smelling salts. Smelling salts produce ammonia gas slowly at room temperature. This gas acts on the brain and prevents fainting or dizziness. In other words, it arouses consciousness.
      6. Ammonia is used in the manufacture of dyes, wood pulp, plastics and fibres such as nylon.

                                              Self-evaluation Test 2.4

      1. What precaution would you take when preparing ammonia in the laboratory?
      2. Calcium oxide is preferred in drying ammonia gas over other drying agents.Explain.
      3. Uwimana, is a senior 3 student. She wanted to confirm whether a gas in a gas jar was ammonia. What chemical test would you advise her to carry out?
      4. Draw the apparatus used to prepare aqueous ammonia in the laboratory.
      5. Large-scale preparation of ammonia is carried out through the Haber process.Name two raw materials and two conditions required for this process.

      6. Ammonia (NH3 ) is a very useful chemical but for many years people could not make large amount of it. In 1908, a chemist called Fritz Haber invented the process for manufacturing ammonia on a large scale. That process is called the Haber process after its inventor. If nitrogen and hydrogen are mixed together with the right conditions, they will react to form ammonia. The chemical reaction has to be carried out at 450°C with a pressure 200 atmospheres air pressure. Finally, the reaction needs a catalyst (iron) to make it work. Most of the nitrogen and hydrogen gas are not used at first so they are channelled back into the reaction vessel until they are all used up. The chemical equation for the reaction is:

      N2 (g) + 3H 2(g) 2NH3 (g)
      Use the write up above to answer these questions
      (a) Ammonia is made out of many atoms.
      (i) How many hydrogen atoms are there in ammonia molecule?
      (ii) How many nitrogen atoms are there in an ammonia molecule?
      (b) What is the most common use of ammonia?
      (c) Who invented the Haber process?
      (d) What does the Haber process produce?
      (e) What type of catalyst is used in the Haber process?
      (f) What happens to the unreacted hydrogen and nitrogen after the ammonia is removed.
      (g) What is the ratio of nitrogen to hydrogen by volume entering the catalyst chamber?
      A. 3 volumes of H2 : 3 volumes of N2
      .B. 1 volumes of H2 : 3 volumes of N2
      C. 3 volumes of H2 : 1 volumes of N2
      D. 1 volumes of H2 : 1 volumes of N

      c) Nitric acid

      Laboratory preparation of nitric acid

      Activity 2.12

      Apparatus and reagents

      Retort glass bottle, round or flat-bottomed flask, concentrated sulphuric acid, potassium nitrate solution, glass cork, clamp stand, source of heat, running water from a tap

      .Procedure

      1. Arrange the apparatus as shown below.


      2. Put some potassium nitrate in the retort glass bottle. Add concentrated sulphuric acid to cover it.
      3. Open the tap water to cool the round or flat bottomed flask.
      4. Heat the retort glass bottle gently.
      5. Note down your observations in the retort glass bottle and flask
      .6. Keep the acid formed for the next activity.

      Study questions

      1. Why are all the apparatus made of glass in this experiment?
      2. What is the importance of tap water?

                           Discussion corner!
      • Discuss the observations made in your study group
      .• Compare your findings with those of other groups.

      I have discovered that...
      When the mixture is heated, effervescence is seen and red-brown fumes are seen in the retort flask.

      The facts

      Potassium nitrate reacts with sulphuric acid to form potassium hydrogen sulphate and nitric acid.

      Potassium nitrate + sulphuric acid potassium hydrogen sulphate + nitric acid
      KNO3(s) + H2SO4 (l) KHSO4 (aq) + HNO3 (aq)

      Nitric acid is more volatile than concentrated sulphuric acid and hence it vaporises.These vapours condense to form nitric acid (a yellow liquid).The red-brown vapours seen in the retort flask are as a result of the decomposition of some of the nitric acid by heat, forming red-brown nitrogen dioxide.

      Nitric acid water + nitrogen dioxide + oxygen
      4HNO3(aq) 2H2O(l) + 4NO2(g) + O2(g)

      The nitric acid collected in this experiment is red-brown because it contains some dissolved nitrogen dioxide, otherwise, pure nitric acid is colourless. The red-brown colour can be removed by bubbling air through the acid.The apparatus used in this experiment is entirely made of glass because hot nitric acid vapours is highly corrosive and would attack cork or rubber stoppers and tubings.

      Potassium nitrate is preferred in preparing nitric acid in the laboratory because it has no water of crystallisation which can dilute the acid produced. Sodium nitrate can also be used but it is hygroscopic.

      Sodium nitrate + sulphuric acid sodium hydrogen sulphate + nitric acid
      NaNO3(s)+ H2SO4 (aq) NaHSO4 (aq)+ HNO3 (aq)

      Note: Nitric acid is kept in dark bottles. This is because light decomposes the acid into water, nitrogen dioxide and oxygen.

      Further activity

      In groups, discuss the industrial manufacture of nitric acid by Ostwald process.Prepare a report on your findings and present it to the class.

      Manufacture of nitric acid by Ostwald Process
      This process involves the catalytic oxidation of ammonia.
      Raw materials
      The raw materials required include
      :• Ammonia-obtained through the Haber process.
      • Air and water.

      The optimum conditions needed are:

      • A temperature of between 850°C - 900°C
      • Platinum - rhodium catalyst
      Note: The air and ammonia must be free of dust and other impurities to avoid poisoning of the catalyst.

      Reactions that take

      place In the first stage, reaction in this process produces nitrogen monoxide and steam. The heat produced in this step maintains the catalyst temperature and therefore no further heating is required.

      Ammonia + oxygen nitrogen monoxide + steam
      4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g)

      The nitrogen monoxide and steam are cooled by air from the compressor. In the second stage, nitrogen monoxide is further oxidised by the air to produce nitrogen dioxide.

      Nitrogen monoxide + oxygen nitrogen dioxide
      2NO (g) + O2 (g) 2NO2 (g)

      In the third stage, nitrogen dioxide is passed through water to form nitric acid (HNO3) and nitrous acid (HNO2).
      Nitrogen dioxide + water nitrous acid + nitric acid
      2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)

      Excess air is passed to oxidise nitrous acid to nitric acid.
      Nitrous acid + oxygen nitric acid
      2HNO2 (aq) + O2 (g) 2HNO3 (aq)



      The nitric acid produced by this process is about 65 % pure. However, the purity of the acid can be increased by distillation. The unreacted air, ammonia and nitrogen dioxide are re-circulated back to the compressor.

      Properties of dilute nitric acid

      Activity 2.13

      Apparatus and reagents

      Test tubes, wooden splint, sodium carbonate, sodium hydroxide, copper(II) oxide, magnesium ribbon, dilute nitric acid, phenolphthalein indicator, zinc metal.

      Procedure

      1. Copy Table 2.4 in your notebook.
      2. Carry out procedures “a-d” and use your results to fill the table.

      Table 2.4: Investigating the properties of dilute nitric acid

      Study question

      Write chemical equations for the reactions that take place in procedure a, b, c and d.

      Discussion corner!

      • In your study groups, discuss the results of the experiment and the study question above.
      • Write a report and submit to the teacher for evaluation.

      i have discovered that...

      There is effervescence in reaction of dilute nitric acid with magnesium ribbon and zinc metal. The magnesium ribbon finally disapears and a colourless solution is formed. Sodium hydroxide solution is neutralised by dilute nitric acid in procedure (a).

      The facts

      Dilute nitric acid acts as any other mineral acid on carbonates, hydroxides and some oxides. The following are chemical properties of dilute nitric acid.

      a) Nitric acid reacts with metal carbonates and hydrogen carbonates to form salt,water and carbon dioxide gas.

      Metal carbonate + dilute nitric acid metal nitrate + water + carbon dioxide
      For example
      Calcium carbonate + nitric acid calcium nitrate + water + carbon dioxide
      CaCO3 (s) + 2HNO3 (aq) Ca(NO3)2 (aq) + H2O (l) + CO2 (g)

      b) It reacts with metal hydroxides to form salt and water. This is a neutralisation reaction.

      Metal hydroxide + dilute nitric acid metal nitrate + water.
      For example:
      Sodium hydroxide + nitric acid sodium nitrate + water
      NaOH (aq) + HNO3 (aq) NaNO3 (aq) + H2O (l)

      c) It reacts with metal oxides to form salt and water.
      Metal oxide + dilute nitric acid metal nitrate + water ,for example:

      Copper (II) oxide + nitric acid copper (II) nitrate + water
      CuO(s) + 2HNO3 (aq) Cu (NO3)2(aq) + H2O (l)

      Nitric acid has oxidising properties that is, it donates its oxygen very easily). It therefore does not usually give off hydrogen with metals. This is because the hydrogen produced is immediately oxidised to form water.For example, zinc metal reacts with dilute nitric acid to form zinc nitrate, nitrogen dioxide and water.

      Zinc + nitric acid zinc nitrate + nitrogen dioxide + water
      Zn(s) + 4HNO3 (aq) Zn(NO3)2 (aq) + 2NO2 (aq) + 2H2O(l)

      An exception is the reaction of magnesium with very dilute and cold nitric acid.
      Magnesium + nitric acid magnesium nitrate + hydrogen
      Mg(s) + 2HNO3 (aq) Mg (NO3)2 (aq) + H2(g)

      Properties of concentrated nitric acid

      Concentrated nitric acid is a very powerful oxidising agent. Its oxidising power depends on the concentration, temperature and the type of reaction.

      Activity 2.14

      Apparatus and reagents

      Iron (II) sulphate solution, dilute sulphuric acid, sulphur powder, 50% concentrated nitric acid, distilled water, barium chloride or barium nitrate solution, lead (II) nitrate solution, copper turnings.

      Procedure
      1. Copy Table 2.5 in your notebook.
      2. Carry out the procedures outlined in the table below and record your observations and conclusions in the blank spaces.

      Table 2.5: Properties of concentrated nitric acid

      Study questions

      1. Which reactions show the oxidation property of concentrated nitric acid?

                                  Discussion corner!
      • In your study groups, discuss the observations made in the procedures above and the study questions
      .• Draw conclusions and write a report.

      I have discovered that...
      Concentrated nitric acid is a strong oxidising agent. In all reactions, it acts as an oxidizing agent. Concentrated nitric acid is reduced to either nitrogen monoxide or nitrogen dioxide.

      The facts

      a) Reaction of concentrated nitric acid with iron (II) sulphate

      Green iron (II) sulphate solution acidified with dilute sulphuric acid reacts with concentrated nitric acid to give a yellow-brown solution of iron (III) sulphate.Concentrated nitric acid therefore oxidises iron (II) ions to iron (III) ions. In the process, the acid is reduced to nitrogen monoxide gas.
      Fe2+ (aq) Fe3+(aq) + e– (oxidation)

      The same type of reaction can take place with iron (II) chloride solution. However,in this case, addition of sulphuric acid is not necessary.
      The nitrogen monoxide produced is readily oxidised to red-brown nitrogen dioxide gas by oxygen from the air.

      Nitrogen monoxide + oxygen nitrogen dioxide
      2NO (g) + O2 (g) 2NO2 (g)
      colourless                       red-brown

      b) Reaction of concentrated nitric acid with copper metal

      Copper metal reacts with 50% concentrated nitric acid producing copper (II) nitrate, water and nitrogen monoxide

      Copper + 50% conc. nitric acid copper (II) nitrate + water + nitrogen monoxide
      3Cu(s) + 8HNO3 (aq) 3Cu (NO3)2 (aq) + 4H2O (l) + 2NO (g)

      The nitrogen monoxide is oxidised to nitrogen dioxide by oxygen from the air.Concentrated nitric acid reacts with copper to form copper (II) nitrate, water and nitrogen dioxide.

      Copper + nitric acid copper (II) nitrate + water + nitrogen dioxide
      Cu(s) + 4HNO3 (l) Cu(NO3)2(aq) + 2H2O (l) + 2NO2(g)

      The same type of reaction takes place when concentrated nitric acid reacts with zinc.

      Zinc + nitric acid zinc nitrate + water + nitrogen dioxide
      Zn(s) + 4HNO3 (l) Zn(NO3)2(aq) + 2H2O (l) + 2NO2(g)

      c) Reaction of concentrated nitric acid with sulphur

      Concentrated nitric acid reacts with sulphur to form sulphuric acid, water and nitrogen dioxide.

      Sulphur + nitric acid sulphuric acid + water + nitrogen dioxide
      S(s) + 6HNO3 (aq) H2SO4 (aq) + 2H2O (l) + 6NO2 (g)

      When barium chloride or barium nitrate solution is added to the sulphuric acid produced in the reaction, a white precipitate of barium sulphate is formed. The sulphuric acid will also form a white precipitate of lead (II) sulphate when reacted  with lead (II) nitrate solution. This is the test for sulphate ions (SO42–). However,if sulphite ions (SO32–) were present, the precipitate would have dissolved in dilute nitric acid.

      The ionic equations for the reactions are:
      Ba2+ (aq) + SO42– (aq) BaSO4(s)
      Pb2+ (aq) + SO42– (aq) PbSO4(s)

      Uses of nitric acid

      1. Manufacture of nitrogenous fertilisers such as sodium nitrate, potassium nitrate, ammonium nitrate and calcium ammonium nitrate.
      2. Manufacture of explosives like trinitrotoluene (TNT) and dynamite.
      3. Manufacture of dyes.
      4. Manufacture of plastics.
      5. It is used as an oxidising agent in textile industries.
      6. It is used in refining of gemstones.

      Self-evaluation Test 2.5

      1. Write two chemical equations for reactions in which nitric acid acts as an oxidising agent.
      2. Nitric acid is not only an important reagent in the Chemistry laboratory but finds application in other fields. Elaborate.
      3. A Senior 3 student added concentrated sulphuric acid to potassium nitrate crystals.
      (a) What observation did the student make?
      (b) Write a balanced chemical equation for the reaction that took place.
      (c) Why is potassium nitrate preferred in this reaction over other nitrates?
      (d) Sodium nitrate can also be used in this reaction. Why is it not preferred?
      4. Give two raw materials and two conditions required in Ostwald process.
      5. Which catalyst is used in large scale preparation of nitric acid?
      A. Vanadium (V) oxide
      B. Pure platinum
      C. Iron
      D. A platinum-Rhodium alloy (Pt - Rh)
      6. Why is air filtered to remove dust and oil before being used in the manufacture of ammonia?
      A. They may poison the catalyst.
      B. They will colour the nitric acid.
      C. They will react with the nitric acid.
      D. They will block the filters.
      7. What word can be used to describe the reaction during absorption of ammonia into water?
      8. Identify the most common fertiliser made from nitric acid.
      A. Nitrogen
      B. Ammonium sulphate
      C. Ammonium nitrate
      D. Manure
      9. Write the equation for the action of heat on potassium nitrate.
      10. Name the nitrate which on heating gives oxygen as the only gaseous product.
      11. Describe what you will observe when concentrated nitric acid is added to copper.
      12. Write the equation for the preparation of nitric acid from potassium nitrate.

      Environmental effects during the manufacture of nitric acid
                                     
                                               Health check
      • Concentrated nitric acid and its vapours are highly corrosive to the eyes, skin and mucous membranes. Dilute solutions of the acid on the other hand cause mild skin irritation and hardening of the epidermis

      .• Contact with concentrated nitric acid burns the skin and produces deep painful wounds. Eye contact can cause severe burns and permanent eye damage. Inhalations of high concentrations can lead to severe respiratory irritations and delayed effects including pulmonary oedema which may be fatal.

      • Ingestion of nitric acid may result in burning and corrosion of the mouth, throat and stomach. An oral dose of 10 ml can be fatal in human beings.

                                 Always be careful when handling nitric acid.

      c) Nitrates

      Nitrates are salts formed when the hydrogen ions in nitric acid are replaced by a metal or ammonium ion

      .Laboratory preparation of nitrates

      Activity 2.15

      Apparatus and reagents

      Spatula, Bunsen burner, filter funnel, filter papers, beakers, stirring rod, measuring cylinder, 1M sodium hydroxide solution, magnesium ribbon, calcium carbonate,1M nitric acid.

      Procedure

      1. Measure 25 cm3 of 1M nitric acid solution and place in a beaker. To this, add 25 cm3 of 1M sodium hydroxide solution and stir. Heat the resulting solution to dryness.
      • Note down your observations.
      2. Add small pieces of magnesium ribbon to 25 cm3 of dilute nitric acid solution in a beaker. Filter the solution to remove excess magnesium and heat the resulting solution to dryness.
      • Note down your observations.
      3. Add excess calcium carbonate to 25 cm3 of nitric acid in a beaker. When the reaction is complete, filter the mixture and heat the filtrate to dryness.
      • Note down your observations.

      Study questions

      1. Why is it important to add excess magnesium ribbons and calcium carbonate in procedures 2 and 3 respectively?
      2. What is the colour of the salt formed in each procedure above?

                   Discussion corner!
      • Discuss the observations made in your study group.
      • Present your finding to the class.

      I have discovered that...
      There are three methods used to prepare nitrates in the laboratory. In all these cases, white solid crystals are formed.

      The facts

      The three methods used to prepare nitrates are:

      (i) Neutralisation of dilute nitric acid by a base/alkali.

      For example:
      HNO3 (aq) + NaOH (aq) NaNO3(aq) + H2O(l)

      When the resulting solution is heated to dryness, a white solid nitrate is obtained.This method is suitable for preparation of nitrates of potassium, sodium and ammonium.(ii) Reaction

      (ii) Reaction between dilute nitric acid and a metal carbonate.

      Examples of such reactions include:

      2HNO3 (aq) + CaCO3(s) Ca(NO3)2(aq) + CO2 (g) + H2O(l)
      2HNO3 (aq)+ PbCO3(s) Pb(NO3)2(aq) + CO2(g) + H2O(l)

      The solution formed is then evaporated to concentrate it and allow crystallisation to take place during cooling. The mother liquor is poured off and the crystals dried between filter papers.

      (iii) Action of nitric acid on a metal.

      This method is not suitable for use with very reactive metals like sodium and potassium. It is also not suitable for preparing salts that are readily oxidised. An example of a salt that can be prepared by this method is magnesium nitrate.

      Mg(s) + 2HNO3 (aq) Mg (NO3)2(aq) + H2 (g)

      Properties of nitrates

      Action of heat on nitrates

      Activity 2.16


      Apparatus reagents

      5 test tubes, labels, blue litmus paper, wooden splint, sodium nitrate, potassium nitrate, copper (II) nitrate, lead(II) nitrate, silver nitrate (if available).

      Procedure

      1. Copy Table 2.6 in your notebook. 

      Table 2.6: Action of heat on metal carbonates


      2. Put a spatulaful of each of the metal nitrates provided in different test tubes and label them accordingly.
      3. Heat each nitrate.
      4. Test any gas evolved using a glowing splint and moist blue litmus paper.
      5. Observe the residue both when hot and cold.
      6. Record all observations in a table.

      Study questions

      1. What is the colour of the gas or the gas mixtures produced in each case?
      2. What is the colour of the hot and cold solid residues in each case?

                                              Discussion corner!
      • In your study groups, discuss the results obtained in the experiment and the study questions above.
      • Prepare a report on how nitrates are affected by heat and present it to the teacher for evaluation.

      I have discovered that...
      The ease of decomposition of metal nitrates increases with decrease in reactivity of the metal constituting the nitrate. The nitrates decomposed form a red-brown acidic gas and other reaction produce a gas that relights a glowing splint.

      The facts

      All nitrates decompose on heating.
      (i) Sodium and potassium nitrates decompose on heating to form metal nitrite and oxygen gas.

      KNO3(s) 2KNO2(s) + O2 (g)
      2NaNO3(s) 2NaNO2(s) + O2 (g)

      (ii) Nitrates of moderately reactive metals like copper, lead and zinc decompose to form metal oxides, nitrogen dioxide and oxygen gas.

      2Pb(NO3)2(s) 2PbO(s) + 4NO2(g) + O2(g)
      2Zn(NO3)2(s) 2ZnO(s) + 4NO2(g) + O2(g)
      2Cu(NO3)2(s) 2CuO(s) + 4NO2(g)+ O2(g)

      (iii) Nitrates of least reactive metals like silver and mercury decompose to form the respective metal, nitrogen dioxide and oxygen gas.

      2AgNO3(s) 2Ag(s) + 2NO2 (g) + O2(g)
      Hg (NO3)2(s) Hg(l) + 2NO2 (g) + O2(g)

      Fumes of mercury are poisonous and so the reaction should not be carried out under normal school laboratory conditions.
      Lithium nitrate behaves differently from other alkali metals as follows.
      4LiNO3(s) 2LiO(s) + 4NO2 (g) + O2 (g)

      Test for nitrates

      Activity 2.17

      The Brown-ring test for nitrates

      Apparatus and reagents

      Test tubes, metal nitrate solution, freshly prepared iron (II) sulphate solution,concentrated sulphuric acid

      Procedure

      1. In a test-tube, mix equal volumes of a metal nitrate solution with freshly prepared iron (II) sulphate solution.
      2. With the test-tube tilted at an angle, carefully pour concentrated sulphuric

      acid down its wall so that two layers of solutions are formed.

      3. Record your observation.
      4. Shake the solution and record what happens.

      Study questions

      1. Why should we use freshly prepared iron (II) sulphate solution?
      2. Explain what happened when you shook the test tube containing the mixture.

                       Discussion corner
      !• Discuss in your study group the observations above.
      • Compare your findings to those of other class member.

      I have discovered that...
      .A brown ring forms between sulphuric acid and the rest of the solution. However, the ring disappears when the solution is shake

      The facts

      Concentrated sulphuric acid reacts with any nitrate salt to produce nitric acid. The nitric acid is reduced to nitrogen monoxide by iron (II) sulphate. The nitrogen monoxide formed reacts with unreacted iron (II) sulphate to form a brown nitroso-iron (II) sulphate (FeSO4.NO) which appears as a brown ring. This is the test for nitrate ions presence in solution.

      Iron(II) sulphate + nitrogen monoxide nitroso-iron(II) sulphate
      FeSO4 (aq) + NO (g) FeSO4.NO (aq)

      The ring disappears if the solution is shaken. This is because concentrated sulphuric acid mixes with the aqueous solution producing a lot of heat. This heat decomposes nitroso-iron(II) sulphate as shown in the equation below.
      Nitroso-iron (II) sulphate iron (II) sulphate + nitrogen (II) oxide
      FeSO4.NO(aq) FeSO4(aq) + NO(g)

      Iron (II) sulphate solution is oxidised by oxygen from the air to iron(III) sulphate within a short time. This is the reason why it should always be freshly prepared.Note: All nitrate salts are soluble in water. They all form nitrate ions when in solution which gives the “brown ring”.

      Uses of nitrates

      1. They are used as fertilisers for example sodium nitrate, potassium nitrate,ammonium nitrate, calcium ammonium nitrate and NPK fertilizers
      .2. Potassium nitrate is used in the manufacture of gun powder. (Gun powder is a mixture of potassium nitrate, charcoal and sulphur). When ignited, the mixture burns and explodes.
      3. Ammonium nitrate is used as a fertiliser and also in the manufacture of explosives.
      4. Nitrates in form of nitroglycerin are used in medicine for treatment of heart diseases.


      Further activity

      Research and make a presentation on the uses of different compounds of nitrogen.

      2.4 Environmental issues with nitrogen dioxide and nitrate fertilisers

      Research activity

      1. In groups, find out the environmental issues associated with nitrogen oxides and nitrate fertilisers.
      2. Suggest ways that can be used to prevent pollution by nitrogenous compounds.
      3. Write a report and hand it to the teacher for evaluation.

      (a) Nitrogen dioxide

      Nitrogen dioxide is found in the atmosphere as a result of industrial processes and emissions from automobile exhausts.Nitrogen dioxide is toxic and can cause respiratory diseases.Nitrogen dioxide also reacts with water vapour

      in the atmosphere and/or rain water to form dilute nitric acid. Rain containing this acid or other acids is called “acid rain”. Acid rain may interfere with animals and plants life on land and in water bodies. This is because the acid lowers the pH of soil and water making it difficult for some organisms to survive.Acid rain also attacks metallic objects such as iron roofs corroding them. This causes leaking in roofs.

      (b) Nitrate fertilisers

      The continued use of large quantities of nitrate fertilisers over long period of time pollutes the environment. This is because excess nitrates in the soil are washed down into rivers, lakes and other water bodies.Once in water, nitrates cause aquatic plants (e.g. plankton and algae) to grow in large quantities (a process known eutrophication). As a result dead plants accumulate faster than can be decomposed by bacteria and they settle to the bottom of the water bodies. Dead plant matter has a high biological oxygen demand (BOD) and some water bodies give up nearly all their dissolved oxygen in order to allow this natural decay process (oxidation) of the dead plant material to take place. This causes a decrease in the concentration of dissolved oxygen in water and this eventually leads to death of aquatic animals such as fish and plants.Also due to large masses of dead plants and animals, the penetration of light in water decreases thereby affecting photosynthesis in aquatic plants.If nitrates are consumed in drinking water, they are converted by bacteria in the intestines to more toxic nitrites. After absorption into the blood stream, the nitrites combine with haemoglobin forming metahaemoglobin. As a result, the oxygen carrying capacity of blood is reduced. Babies suffering from excess intake of nitrates have been known to change ther colour to “blue” due to lack of oxygen in their blood, a condition referred to as ‘blue babies’.

                                             
                                                 My environment, my life!
      We should advise people in our community to avoid cultivating near river banks. This will prevent the washing away of fertilisers into water bodies hence prevent eutrophication.

      Unit summary 2

      • Nitrogen can be isolated from air by fractional distillation of liquid air. The process involves removing other components of air like carbon dioxide,oxygen, noble gases and moisture (water vapour).

      • In the laboratory, nitrogen is prepared by reacting ammonium chloride and sodium nitrite.

      • Nitrogen is inert at room temperature because of the strong triple covalent bonds between the two nitrogen atoms which make the molecule.

      • In the laboratory, nitrogen dioxide is prepared by reacting copper (II) nitrate with concentrated nitric acid.
      • Nitrogen dioxide is a red-brown gas with an irritating pungent smell and dissolves in water to form nitrous and nitric acids. It is denser than air.

      • Ammonia is an alkaline gas prepared in the laboratory by reacting calcium hydroxide and ammonium chloride. The flask containing reacting reagents should be slanting to prevent moisture that forms in the cooler parts of the apparatus from flowing back into the flask which can can cause it to crack.

      • Large scale manufacture of ammonia is done through Haber process. The raw materials required are nitrogen and hydrogen in the ratio 1:3. Other conditions required are pressure of 200 - 250 atmospheres and a temperature of 4000C - 5000C.

      • Ammonia is used in the manufacture of fertilisers, nitric acid, dyes, wood pulps, plastics and fibres.

      • Laboratory preparation of nitric acid is done by reacting potassium nitrate and sulphuric acid.

      • Nitric acid is prepared on large scale through Ostwald process. The raw materials needed are ammonia, air and water.

      • Concentrated nitric acid is a strong oxidising agent. It is reduced to either nitrogen monoxide or nitrogen dioxide.

      • Nitric acid is used in the manufacture of nitrogenous fertilisers, explosives,dyes and plastics. It is also used as an oxidizing agent in textile industries.

      • Nitrates are salts formed when the hydrogen ion in nitric acid is replaced by a metal or an ammonium ion.

      • The three methods used to prepare nitrates are neutralisation of dilute nitric acid by an alkali, reaction between dilute nitric acid and a metal carbonate and action of nitric acid on a metal.

      • All nitrates decompose on heating. However, the degree of decomposition increases with decrease in reactivity of the metal forming the nitrate.94Test your Competence 2

      • Brown ring test is a test used to identify nitrates. It involves mixing equal volumes of a nitrate with freshly prepared iron (II) sulphate. Concentrated sulphuric acid is then carefully added.

      • Nitrogen dioxide released into the atmosphere pollutes the environment. It dissolves in rain water forming acid rain. Acid rain destroys vegetation, animals and buildings. Excess use of nitrate fertilisers also leads to water pollution and eutrophication.

      Test your Competence 2

      1. Cultivating near water bodies leads to environmental pollution. Explain this with reference to effects of nitrogen compounds on the environment.

      2. Nitrogen dioxide is a _________ gas in colour that has a _________ smell while ammonia is a colourless gas with a _________ smell.
      3. Mutesi dissolved nitrogen dioxide gas in water. She later put blue and red litmus papers in the solution formed. The blue litmus paper turned red while there was no effect on the red litmus paper.
      (a) Explain the observations made above.
      (b) Write a chemical equation for the reaction that took place.
      4. In the preparation of ammonia, a student set up an experiment as shown below.


      (a) Identify and correct the mistake in the set up.
      (b) Name substance A and state its function.
      (c) Name the method used to collect ammonia gas in the set up and explain why it is suitable.
      (d) Write a chemical equation for the reaction that took place in the round- bottomed flask.
      5. Large scale manufacture of ammonia is carried out through ______ process while manufacture of nitric acid in large scale is done through ______ process.
      6. Name two raw materials and two conditions necessary for the large scale manufacture of ammonia.
      7. Study the flowchart below then answer the questions that follow.


      (a) Name the process illustrated in the flowchart above ––––––––––––––.
      (b) Identify (i) Gas X _________(ii) Product Z _________
      (c) Write a chemical equation for the reaction that takes place in the catalytic chamber
      . (d) Why is it recommended that all the apparatus in the laboratory preparation of this product be made of glass?
      8. The following methods are suitable for preparation of nitrates except _______?
      A. Reacting dilute nitric acid with sodium hydroxide.
      B. Reacting sodium metal with nitric acid.
      C. Reacting nitric acid with lead carbonate.
      D. Reacting dilute nitric acid with magnesium.
      9. Write true or false for each of the following:
      (a) All nitrates decompose on heating. _________
      (b) Ammonia is the only common basic gas. _________
      (c) Nitrogen dioxide dissolves in water to form nitric acid only. _________
      (d) Brown-ring test is used to identify nitrates.
      • Key unit competency

        After studying this unit, I should be able to relate the properties of sulphur and its compounds to their uses, describe how some compounds of sulphur are prepared and discuss the related environmental issues.

        Learning objectives

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

        • Recall the occurrence, extraction, properties and uses of sulphur.

        • Prepare, test and collect sulphur dioxide gas and explain the impact of sulphur oxides on the environment.

        • Describe the industrial preparation of sulphuric acid by the contact process.

        • Develop skills in observation in preparing sulphur dioxide gas and testing for the presence of sulphates and sulphites in given solutions.

        • Protect natural resources.

        • Develop self confidence in discussions and presentation of research findings.

        • Develop a culture of working in a team during research and discussions.

        • Respect procedures during experiments.

        Mind teaser

        Study the chart below carefully. What is the chart about?