Carbon and Its Compounds - High School Chemistry

- All living things contain carbon compounds.
- These include carbohydrates, proteins, fats among others. Petrol, oil and wood are carbon compounds.
- Carbon is incorporated in plants from carbon(IV) oxide during photosynthesis.
- Carbon is one of the most useful and important elements.
- It is located in period two, group IV of the periodic table. Figure 1

Figure 1: Position of carbon in the periodic table

Carbon atom has atomic number 6 and electron arrangement of 2.4

Figure 2: Atomic structure of carbon

- There are four electrons in the outermost energy level in a carbon atoms.
- It is very difficult for a carbon atom to either lose or gain electrons.
- Therefore, carbon can only bond by sharing of electrons to gain stability.

- Carbon occurs in the pure state in two allotropic forms; diamond and graphite.
- Allotropy is the existence of an element in more than one form but in the same state.
- It also occurs in many impure forms (charcoal).
- Charcoal is also made up of the graphite lattice.

a) Diamond

Structure of Diamond

- In diamond, each carbon atom is joined to four other carbon atoms.
- Each carbon atom shares electrons with each of its neighbours forming strong covalent bonds.
- Therefore all 4 valence electrons are involved in bonding, See Figure 3

Figure 3: Carbon atom bonded to four other carbon atoms

- In diamond, a giant atomic structure is formed by three dimensional structure.
- Each carbon atom is surrounded tetrahedrally by four others.
- Each carbon atom is bonded by very strong covalent bond to the other carbon atom.

Figure 4: Diamond giant atomic structure

Properties of Diamond

- Diamond is very hard.
- Carbon atoms in diamond are linked by very strong covalent bonds.
- The atoms in diamond are not arranged in layers, therefore they cannot slide over one another like the atoms in metals or graphite.
- This makes diamond the hardest substance ever known.
- Diamond has a very high melting and boiling point.
- Since diamond is held by very strong covalent bonds, the atoms cannot vibrate enough to break away the bonds.
- This results in diamond having very high melting and boiling points.
- Diamond does not conduct electricity.
- All valency electrons in diamond are involved in bonding.
- There are no delocalised electrons to conduct electricity like in metals or graphite.
- Diamond has a very shiny sparkling appearance.

Uses of Diamond

1. It is used for making jewelry e.g. necklaces, rings and earrings.
2. Diamond is used to cut glass and in drilling equipment because it is very hard.

Structure of Graphite

- In graphite, one carbon atom forms covalent bonds with three others to give a ring of six carbon atoms.
- Since there are four electrons available for bonding in a carbon atom, one electron is left free.
- When an electron is free to move within the structure, it is said to be delocalised.
- Many of these hexagonal rings join together by strong covalent bonds to form a layer so that every layer has a giant atomic structure.
- These giant atomic structures make flat sheets that lie on top of each other held together by weak van der Waals forces.

Figure 5: Structure of graphite showing the layers

Properties of Graphite

1. Soft and slippery
- Graphite is soft and slippery because the sheets of atoms can slide over each other easily due to the weak Van der waals forces between the parallel sheets of giant atomic structures.
2. Good conductor of electricity
- This is because each carbon atom has four outer electrons but forms only three covalent bonds.
- The fourth electron is free to move through the graphite, carrying a charge.
- Therefore there are many delocalised electrons to conduct an electric current.
3. Graphite has a very high melting
- Although the layers of graphite move over each other easily, it is difficult to break the bonds between carbon atoms within one layer.
- Covalent bonds are strong, hence require a lot of energy to break. Due to this, graphite melts at 3,370^°C and boils at 4,830^°C.

Uses of Graphite

1. Graphite is used as lubricant because it is soft and slippery.
2. It is used to reinforce metals and broken bones.
3. Graphite is used as the positive terminals in dry cells and as electrodes in industries.

c) Charcoal

Structure of Carbon

- Charcoal is an amorphous (without shape) form of carbon and exists in many forms.
These include:
a) Animal charcoal - formed when bones are heated in limited supply of air.
b) Wood charcoal - which is the solid left when wood is heated in limited supply of air.
c) Sugar charcoal - is formed by dehydrating cane-sugar or glucose with concentrated sulphuric acid or heating the sugar in the absence of air.
d) Lamp black - is formed when petroleum, kerosene, turpentine, natural gas and other hydrocarbons burn in a limited supply of air.
e) Coke - is the solid left when coal is heated in absence of air.
f) Soot - is formed when there is incomplete combustion of fuels. It is found in chimneys of houses, lantern lamps, and many other places.

Uses of Carbon

1. Because of its good absorption power, amorphous carbon is used to purify sugar in industries.
2. It can also be used in gas masks for people who work in industries which produce poisonous gases since it absorb gases.

A summary of Properties of Allotropes of Carbon

	Diamond	Graphite	UAmorphous Carbon
Appearance	Colourless, transparent crystals that sparkle in light	Dark-grey, opaque and shining	Black, opaque and very dull
Density	Highest (3.5g/cm3)	Moderate (2.3g/cm3)	Low 1.5g/cm3 when air free
Hardness	Hardest natural substance known	Soft and slippery	Soft
Electrical conductivity	Does not conduct electricity	Conducts electricity	Does not conduct electricity

(a) Combustion

- A gas is produced that forms a white precipitate with calcium hydroxide solution.
- All forms of carbon burn in sufficient supply of oxygen to form carbon(IV) oxide and heat.

Carbon + oxygen → carbon(IV) oxide + heat
C(s) + O₂(g) → CO₂(g) + heat

- Diamond and graphite ignite to red-hot with a lot of difficulty.
- If burned in an insufficient supply of air (oxygen), carbon(II) oxide is produced instead.

Carbon + oxygen → carbon(II) oxide + heat
2C(s) + O₂(g) → 2CO(g) + heat

(b) Reducing Nature of Carbon

- In all cases, carbon reduces metal oxides to their respective metals.

Lead(II) oxide + carbon → lead + carbon(IV) oxide
2PbO(s)_(yellow) + C(s) → 2Pb(s)_(silvery) + CO₂(g)

Zinc oxide + carbon → zinc + carbon(IV) oxide
2ZnO(s)_(white) + C(s) → 2Zn(s) + CO₂(g) _(grey)

Iron(III) oxide + carbon → iron + carbon(IV) oxide
2Fe₂O₃(s)_(red-brown) + 3C(s) → 4Fe(s)_(grey) + 3CO₂(g)

Copper(II) oxide + carbon → copper + carbon(IV) oxide
2CuO(s)_(black) + C(s) → 2Cu(s)_(red-brown) + CO₂(g)

(c) Reaction of Carbon with Acids and Steam

- Carbon reduces concentrated sulphuric and nitric acids when heated with them.
Carbon + sulphuric acid → carbon(IV) oxide + sulphur(IV) oxide + water
C(s) + 2H₂SO₄(aq) → CO₂(g) + 2SO₂(g) + 2H₂O(l)

Carbon + nitric acid → carbon(IV) oxide + nitrogen(IV) oxide + water
C(s) + 4HNO₃(aq) → CO₂(g) + 4NO₂(g) + 2H₂O(l)

- These acids donate oxygen to carbon to form carbon(IV) oxide
Carbon + oxygen → carbon(IV) oxide
C(s) + O₂(g) → CO₂(g)

- Water-gas which is a mixture of carbon(II) oxide and hydrogen is produced by passing steam through coke at temperatures higher than 1000^°C.

Carbon + steam → carbon(II) oxide + hydrogen
C(s) _(coke) + H₂O(g)_(steam) → CO(g) + H₂(g)

Carbon(II) oxide are hydrogen and combustible and therefore water-gas is used as a fuel.

Hydrogen + oxygen → steam + heat
2H₂(g) + O₂(g) → 2H₂O(g) + heat

Carbon(II) oxide + oxygen → carbon(IV) oxide + Heat
2CO(g) + O₂(g) → 2CO₂(g) + Heat

Laboratory preparation of carbon(IV) oxide (CO₂)

Figure 6: Preparation of carbon(IV) oxide

- The gas should now be collected by downward delivery
- A steady reaction is observed when marble chips react with dilute hydrochloric acid.

Calcium carbonate + hydrochloric acid → calcium chloride + carbon(IV) oxide + water
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l)

- Ordinary air contains about 0.03 per cent of carbon(IV) oxide.
- Exhaled air contains about 3 per cent.
- The burning of wood, coal, petrol,oils and other carbon compounds add carbon(IV) oxide to the atmosphere.
- Green plants use it to make their food. We also use in soft drinks like Coca-Cola.

Physical Properties of Carbon(IV) Oxide

• Carbon(IV) oxide is a colourless gas.
  
• It has a slight smell.
  
• It is denser than air.
  
• It is slightly soluble in water.
  

Chemical Properties of Carbon (IV) Oxide

a) Reaction of carbon(IV) oxide with calcium hydroxide

- It forms a white precipitate with calcium hydroxide solution.
- This precipitate is called calcium carbonate and it is insoluble in water.

Calcium hydroxide + carbon(IV) oxide → calcium carbonate + water
Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

- When carbon(IV) oxide is passed through the calcium carbonate and water for a long time, the insoluble calcium carbonate is converted into soluble calcium hydrogencarbonate solution.

Calcium carbonate + carbon(IV) oxide + water → calcium hydrogencarbonate
CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq)

b) Reaction of carbon(IV) oxide with water

- Carbon(IV) oxide moderately reacts with water.
- The solution formed is slightly acidic, it does not turn blue litmus paper fully red.
- Some carbonic acid is formed but in very small quantities.

Carbon(IV) oxide + water → carbonic acid
CO₂(g) + H₂O(l) → H₂CO₃(aq)

- Since carbon(IV) oxide reacts with water to form carbonic acid, this acid formed reacts with sodium hydroxide to form sodium carbonate and water only.

Sodium hydroxide + carbon(IV) oxide → sodium carbonate + water
2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l)

- Since carbonic acid is dibasic it forms a normal and an acid salt.
- When carbon(IV) oxide is passed for along time in concentrated solution of sodium hydroxide a further reaction occurs forming sodium hydrogencarbonate, some of which precipitates.

Sodium carbonate + carbon(IV) oxide + water → sodium hydrogencarbonate
Na₂CO₃(aq) + CO₂(g) + H₂O(l) → 2NaHCO₃(aq)

- Potassium hydroxide solution gives a similar reaction but its hydrogen carbonate is more soluble and does not precipitate.
- This is why potassium hydroxide solution is preferred as an absorbent for carbon(IV) oxide.

c) Carbon(IV) oxide and combustion

- Carbon(IV) oxide extinguishes a burning candle showing that it does not support combustion.

d) Reaction of carbon(IV) oxide with magnesium

- The burning magnesium ribbon continues to burn inside a gas jar full of carbon(IV) oxide.
- 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.
- Carbon(IV) oxide does not support combustion but magnesium continues to burn in it. The heat of burning magnesium is sufficient to decompose carbon(IV) oxide into carbon and oxygen.

heat Carbon(IV) oxide → carbon + oxygen
CO₂(g) → C(s) + O₂(g)

- The oxygen produced supports the combustion of magnesium to form magnesium oxide.
Magnesium + oxygen → magnesium oxide
2Mg(s) + O₂(g) → 2MgO(s)

- The general equation can thus be written as follows:
Magnesium + Carbon(IV) oxide → Magnesium oxide + Carbon
Mg(s) + CO₂(g) → MgO(s) + C(s)

- This experiment shows that magnesium displaces oxygen from carbon(IV) oxide.
- Therefore magnesium is more reactive towards oxygen than carbon.

Uses of Carbon (IV) Oxide

1. Carbon(IV) oxide is used in the manufacture of sodium carbonate used for baking cakes, bread among other products and sodium hydrogencarbonate which is used in some health salts such as Eno and Andrews liver salt.
2. Carbonated drinks
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(IV) oxide in water under pressure. It is later sweetened, flavoured and sometimes coloured. The dissolved carbon(IV) gives it a pleasant taste.
3. As a refrigerant
Solid carbon(IV) oxide commonly known as dry ice is preferred to ordinary ice (solid water) because it sublimes at room temperature forming gaseous carbon(IV) oxide and therefore leaves no residue like ordinary ice.
4. Fire extinguishers
Fire extinguishers often contain sodium hydrogencarbonate or sodium carbonate solution and sulphuric acid. When mixed by inversion or pressing a plunge, carbon(IV) oxide is produced due to the reaction of carbonate with sulphuric acid.
5. Making rain during drought or in areas of little rain
In making rain, dry ice (solid carbon(IV) oxide) is spread in the clouds to accelerate the condensation process.
6. In Solvay process
This is a process used in the manufacture of sodium carbonate (soda-ash).
7. It is also used to transfer heat energy from certain types of nuclear reactors.

Note: Preparation and investigation of properties of carbon(II) oxide should not be done practically due to its poisonous nature.

Laboratory Preparation of Carbon (II) Oxide (CO)

Figure 7: Preparation of carbon(II) oxide from concentrated sulphuric acid and sodium methanoate

- Sulphuric acid first liberates methanoic acid from sodium methanoate and then dehydrates it.

Sodium methanoate + sulphuric acid → methanoic acid + sodium hydrogensulphate
HCOONa(s) + H₂SO₄(aq) → HCOOH(aq) + NaHSO₄(aq)

- Potassium hydroxide solution removes any carbon(IV) oxide or sulphur(IV) oxide which may be present.
- The gas is usually collected over water.
- If it is required dry, it may be passed through concentrated sulphuric acid and then collected by upward delivery (being less dense than air).
Note: Methanoic acid can be used instead of sodium methanoate.

Laboratory Preparation of Carbon(II) Oxide from carbon(IV) Oxide

Figure 8: Preparation of carbon(II) oxide from carbon(IV) oxide

Carbon(IV) oxide is reduced by the red-hot charcoal to carbon(II) oxide.
Carbon(IV) oxide + carbon → carbon(II) oxide
CO₂(g) + C(s) → 2CO(g)
- Any unreacted carbon(IV) oxide is removed by potassium hydroxide solution.
Potassium hydroxide + carbon(IV) oxide → potassium carbonate + water
2KOH(aq) + CO₂(g) → K₂CO₃(aq) + H₂O(l)

Physical Properties of Carbon (II) Oxide (CO)

• Colourless gas
  
• Odourless
  
• Tasteless
  
• Less dense than air
  
• Insoluble in water
  

Combustion of Carbon (II) Oxide

- Carbon(II) oxide gas burns with a blue flame forming carbon(IV) oxide gas.
Carbon(II) oxide + oxygen → carbon((IV) oxide
2CO(g) + O₂(g) → 2CO₂(g)
- Carbon(IV) oxide forms a white precipitate of calcium carbonate with calcium hydroxide solution.
- Carbon(II) oxide is said to be a reducing agent because it takes oxygen.
- Oxygen is the oxidizing agent.

- It will be observed that some carbonates decompose when heated and others do not.
- Carbonates of:

- (NH₄)₂CO₃ produces carbon(IV) oxide, water and ammonia gas on heating.
- Therefore, the carbonates of metals high in the reactivity series do not decompose.
- Ease of decomposition of metal carbonates increases as we go down the reactivity series.
- The following are equations showing decomposition of carbonates.

Copper carbonate → copper(II) oxide + carbon(IV) oxide
CuCO₃(s)_(green) → CuO(s)_(black) + CO₂(g)

Magnesium carbonate → magnesium(II) oxide + carbon(IV) oxide
MgCO₃(s)_(white) → MgO(s)_(white) + CO₂(g)

Zinc carbonate → zinc oxide + carbon(IV) oxide
ZnCO₃(s)_(white) → ZnO(s)_(white) + CO₂(g)

Lead carbonate → lead(II) oxide + carbon(IV) oxide
PbCO₃(s)_(white) → PbO(s)_(yellow)) + CO₂(g)

Calcium carbonate → calcium oxide + carbon(IV) oxide
CaCO₃(s)_(white) → CaO(s)_(white) + CO₂(g)

- Copper(II) oxide is black, while oxides of magnesium and calcium are white.
- Zinc oxide is yellow when hot and white on cooling.
- Lead(II) oxide is red-brown when hot and yellow on cooling.
- Sodium carbonate and potassium do not decompose on heating.
- Sodium hydrogencarbonate and calcium hydrogencarbonate decompose on heating to form a carbonate, water and carbon(IV) oxide.

Sodium hydrogencarbonate → sodium carbonate + carbon(IV) oxide + water
2NaHCO₃(s) → Na₂CO₃(s) + CO₂(g) + H₂O(l)

Calcium hydrogencarbonate → calcium carbonate + carbon(IV) oxide + water
Ca(HCO₃)₂(s) → CaCO₃(s) + CO₂(g) + H₂O(l)

- When calcium hydrogencarbonate is heated strongly calcium carbonate formed decomposes to form calcium oxide and carbon(IV) oxide.
- Ammonium carbonate decomposes on heating to form ammonia gas, water and carbon(IV) oxide.

Ammonium carbonate → ammonia + carbon(IV) oxide + water
(NH₄)₂CO₃(s) → 2NH₃ (g) + CO₂(g) + H₂O(l)

- Ammonia and carbon(IV) oxide 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.
Note: Sometimes sodium carbonate may be contaminated with sodium hydrogencarbonate and therefore carbon(IV) oxide can be obtained on heating it.

Action of Dilute Acids on Carbonates and Hydrogencarbonates

- All carbonates liberate carbon(IV) oxide on addition of a dilute acid.
- They also form a salt and water. For example:

Calcium carbonate + nitric acid → calcium Nitrate + water + carbon(IV) oxide
CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)(aq) + H₂O(l) + CO₂(g)

Calcium carbonate + hydrochloric acid → calcium chloride + water + carbon(IV) oxide
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

All hydrogencarbonates also liberate carbon(IV) oxide on addition of an acid. For example:

Sodium hydrogen carbonate + nitric acid → sodium nitrate + carbon(IV) oxide + water
NaHCO₃(s) + HNO₃(aq) → NaNO₃(aq) + CO₂(g) + H₂O(l)

Sodium hydrogencarbonate + hydrochloric acid → sodium chloride + carbon(IV) oxide + water
NaHCO₃(s) + HCl(aq) → NaCl(aq) + CO₂(g) + H₂O(l)

Production and Manufacture of Sodium Carbonate

I. Production of sodium carbonate (soda ash) at Magadi Soda Company

- The most efficient means of trona extraction uses dredging techniques.
- Both a bucket dredge and a cutter-suction dredge are used for this purpose, concentrating on the top 4 metres or so of the solid crust where the purest material is found.
- As the solid material is removed from the surface, liquor (unwanted solution) from the surrounding trona drains into the resulting cavity, forming a pool or paddock in which dredgers can float.
- The crystals of trona produced by the dredgers are mixed with liquor and pumped as slurry back to the ash plant.
- Here the liquor is discarded and the crystal washed and centrifuged.
- The dump crystals are then fed into calciners (kilns).
- Here the residual moisture, water crystallisation and carbon(IV) oxide gas are driven off to leave normal sodium carbonate (soda ash).

- The soda ash from the calciners passes through rotary drum coolers before entering the grinding and screening plant.
- Oversize material is removed and grounded again.
- The final product is conveyed to silos where it is packed or loaded directly into specially bulk raid hopper wagons.
- The liquor from this process is used to extract sodium chloride.
- Solar evaporation process is used. Most of the Magadi sodium chloride is sold within Kenya for livestock consumption or for industrial purposes.

Figure 9: A flow chart of soda ash manufacturing process

II. Manufacture of sodium carbonate (soda ash) by Solvay process

- The main raw materials are sodium chloride and calcium carbonate.
- Concentrated brine (sodium chloride solution) is saturated with ammonia in a tower.
- The ammoniacal brine formed is run down the Solvay tower.
- Carbon(IV) oxide is forced into the tower from the bottom.
- The towers are filled with baffles. These baffles make the liquid flow slowly and increase the surface area for reaction.
- Sodium hydrogen carbonate is formed in Solvay tower. The reaction occurring is as follows.

Sodium chloride + ammonia + carbon(IV) oxide + water → sodium hydrogencarbonate + ammonium chloride
NaCl(g) + NH₃(aq) + CO₂(g) + H₂O(I) → NaHCO₃(s) + NH₄Cl(aq)

- Then sodium hydrogencarbonate is filtered off, and washed to remove ammonium chloride.
- It is then heated in a furnace roaster to give sodium carbonate and carbon(IV) oxide.

Sodium hydrogencarbonate → sodium carbonate + carbon(IV) oxide + water
2NaHCO₃(s) → Na₂CO₃(s) + CO₂(g) + H₂O(l)

- Carbon(IV) oxide formed is recycled to the Solvay tower.
- The other source of carbon(IV) oxide is from heating calcium carbonate(limestone) in the kiln where it dissociates into calcium oxide and carbon(IV) oxide.

Calcium carbonate → calcium oxide + carbon(IV) oxide
CaCO₃(s) → CaO(s) + CO₂(g)

- The calcium oxide formed from above reaction is slaked by addition of water.

Calcium oxide + water → calcium hydroxide
CaO(s) + H₂O(l) → Ca(OH)₂z(s)

Figure 10: Solvay process

- The calcium hydroxide formed is heated with the filtrate, ammonium chloride from the solvay tower, to produce ammonia.
Ammonium chloride + calcium hydroxide → calcium chloride + ammonia + water
2NH₄Cl(aq) + Ca(OH)₂(s) → CaCl₂(aq) + 2NH₃(g) + 2H₂O(l)
- The ammonia from this reaction is returned to the ammoniating tower.
- A careful examination of the flow diagram (Figure 10) shows that the only waste product in this process is calcium chloride.
- The carbon(IV) oxide and ammonia are recycled.
- Therefore, the solvay process is very efficient.
- The raw materials are cheap and readily available and only one waste product is formed.

Uses of sodium carbonate

1. Manufacture of glass.
2. In domestic water-softening.
3. In the manufacture of chemicals e.g. sodium hydroxide.
4. In the manufacture of laundry detergents.
5. In the paper-making process.
6. Textiles.

Carbon cycle

- The processes which add carbon(IV) oxide to the atmosphere are:
1. Combustion:
- Petroleum, coal, wood, charcoal, wax or any other organic compound.
2. Respiration:
- During this process the sugar in the animal bodies is changed by oxygen to carbon (IV) oxide and water.
- Carbon(IV) oxide is breathed out to the atmosphere.
3. Decay of plants and animals:
- Any carbon in plants and animals is converted to carbon(IV) oxide in the decaying process if sufficient oxygen supply is available.
4. Making of beer and wine:
- Yeast changes sugar to ethanol and carbon(IV) oxide. The process is called fermentation.
5. In areas where there are factories manufacturing calcium oxide from calcium carbonate, carbon(IV) oxide is formed as a by-product.
- If not collected for use it finds its way to the atmosphere.
Calcium carbonate → calcium oxide + carbon(IV) oxide
CaCO₃(s) → CaO(s) + CO₂(g)

Carbon(IV) Oxide is removed from the atmosphere by:

1. Water

- Carbon(IV) oxide dissolves in rivers and lakes to form carbonic acid.
- The carbonic acid reacts with calcium carbonate and magnesium carbonate to form calcium hydrogen carbonate and magnesium hydrogencarbonate respectively.
Carbon(IV) oxide + water → carbonic acid
CO₂(g) + H₂O(l) → H₂CO₃(aq)

Calcium carbonate + water + carbon(IV) oxide → calcium hydrogen carbonate
CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)

Magnesium carbonate + Water + Carbon(IV) oxide → Magnesium hydrogen carbonate
MgCO₃(s) + H₂O(l) + CO₂(g) → Mg(HCO₃)₂(aq)

- The hydrogen carbonates cause temporary hardness in water.
- This hard water is used by animals to make shells.

2. Photosynthesis

In this process, green plants use carbon(IV) oxide to make sugar using sunlight as a source of energy.
Carbon(IV) oxide + water → sugar + oxygen
CO₂(g) + H₂O(l) → C₆H₁₂O₆(aq) + O₂(g)

3. Calcium Hydroxide

- Natural calcium hydroxide slowly reacts with carbon(IV) oxide from the air to form calcium carbonate
Calcium(II) hydroxide + carbon(IV)oxide → calcium carbonate + water
Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

Figure 11: Carbon Cycle

Effects of Carbon (IV) Oxide and Carbon (II) Oxide on the Environment

- The Greenhouse effect is the way in which heat is trapped close to Earth's surface by "greenhouse gases" including Carbon (IV) Oxide and Carbon (II) Oxide.
- Generally when we talk about weather, we are talking about temperature, rain or hail storms, typhoons and so on.
- Temperature is the amount of heat found in the atmosphere.
- This heat in the atmosphere comes from the sun. Each day a small amount of the sun's total heat energy reaches the top layers of earth's atmosphere.
- The earth then maintains heat energy balance so that not all the heat coming to the earth remains.
- In time, the heat energy held in the atmosphere and in earth's surface escapes back into space.
- Then earth's atmosphere acts like the glass roof of a greenhouse.
- In a greenhouse the sun's rays pass through the glass windows and warm the plants as seen in Figure 12

Figure 12: A greenhouse

- The sun's rays pass through the space, and onto the soil, rocks and water bodies such as ocean, lakes and rivers.
- The rays are reflected back into space in form of infrared radiations which do not easily escape into atmosphere.
- This is because some gases like carbon(IV) oxide in the lower part of the atmosphere take in some of the heat energy and send it back to earth.
- As a result the temperature of the surface of the earth is raised thereby preventing excessive cooling on the earth.
- This is called the greenhouse effect.
- There is an increasing concern among scientists that large amounts of carbon(IV) oxide in the atmosphere are changing the green house effect.
- This build up of carbon(IV) oxide has resulted from increased burning of coal, oil, natural gases, firewood, etc and extended cutting of trees.
- Trees purify the atmosphere by absorbing carbon(V) oxide and releasing oxygen.
- When a lot of trees are cut down, the carbon(IV) oxide they would have absorbed is released to the atmosphere causing an increase in the earth's temperature.
- This heat has caused melting of ice in parts of the north and south poles and subsequent rise in sea level.
- As a result it has been observed that ocean waters have risen in some parts of the world causing shoreline areas to flood.
- The result is that more areas in the high altitudes can be farming but with more in the low altitude areas being rendered useless for farming because of flooding.
- Also, changing temperatures have destroyed plant and animal life on sea and land, as some organisms cannot survive in areas of extreme temperatures.
- Such changes have also interfered with the world climate causing unexpected heavy rains and droughts in many parts of the world.

Effects of carbon(II) oxide on the environment

- Carbon(II) oxide is a very poisonous gas.
- It has no colour, taste or smell, thus making it hard for anyone to detect its presence.
- It is formed during respiration by incomplete combustion of charcoal, coal, natural (methane) gas, fires, firewood in poorly ventilated areas. and proper ventilation in rooms, industrial areas, etc, helps in preventing the build up of carbon (IV) oxide.

Reducing pollution

- It is important that we reduce the level of air pollution for a healthier environment. - The following are some of the guidelines that will ensure that our environment is kept clean and safe for our future generations.
1. Laws in the cities and other urban areas should restrict the amount of smoke and other waste products released by factories which pollute the air.
- Recycling should be encouraged as much as possible.
2. Burning of rubbish should be restricted such that it is done only in designated areas.
3. Vehicles should have their exhaust pipes fitted with filters to trap impurities and other substances that contaminate the environment.