SSS 1: CARBON AND ITS COMPOUNDS (II)

OXIDES OF CARBON

Carbon forms two important oxides, namely:

1. Carbon(IV) oxide, CO₂

2. Carbon(II) oxide, CO

                                                   Carbon (IV) oxide, CO₂

The atmosphere contains about 0.03% by volume of CO₂. A small percentage of CO₂ is also found dissolved in water.

In the combined form, carbon is found mainly as metallic trioxocarbonates(IV) and hydrogen trioxocarbonates(IV) in the Earth's crust, especially in limestone regions and coral reefs.

Preparation of CO₂

A. Laboratory Preparation

CO₂ is prepared in the laboratory by:

(i) Action of dilute acids on trioxocarbonates(IV) or hydrogen trioxocarbonates(IV)

For example:

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Another example is:

NaHCO₃(aq) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)

(ii) Heating metallic trioxocarbonates(IV)

Except those of sodium and potassium, many metallic trioxocarbonates(IV) decompose on heating.

CuCO₃(s) → CuO(s) + CO₂(g)

However, sodium and potassium hydrogen trioxocarbonates(IV) decompose on heating:

2KHCO₃(s) → K₂CO₃(s) + H₂O(l) + CO₂(g)

                                                    KIPP'S APPARATUS

A Kipp's apparatus is used to provide a supply of CO₂ whenever it is needed in the laboratory.

In the Kipp's apparatus, CO₂ is produced by the action of dilute hydrochloric acid on marble chips (calcium carbonate).

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

B. Industrial Preparation

Industrially, CO₂ is obtained as a by-product of the fermentation process and when limestone is heated to make quicklime.

CaCO₃(s) → CaO(s) + CO₂(g)

Physical Properties of CO₂

1. CO₂ is a colourless and odourless gas.

2. It is about 1.5 times denser than air.

3. It is soluble in water.

CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)

4. It changes damp blue litmus paper to pink.

5. On cooling, it readily liquefies and solidifies at about −78 °C, forming a white solid known as dry ice.

Chemical Properties of CO₂

1. Reaction with Water

CO₂ reacts with water to form weak carbonic acid:

CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)

2. Reaction with Alkalis

CO₂ reacts with alkalis to form trioxocarbonates(IV).

For example:

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

With excess CO₂:

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

3. Reaction with Burning Magnesium

CO₂ does not support ordinary combustion, but burning magnesium can continue to burn in it.

CO₂(g) + 2Mg(s) → 2MgO(s) + C(s)

4. Reaction with Red-Hot Carbon

When CO₂ is passed over red-hot carbon, it is reduced to carbon(II) oxide:

CO₂(g) + C(s) ⇌ 2CO(g)

                                                  Test for CO₂

Limewater, Ca(OH)₂(aq), turns milky when CO₂ is bubbled through it because of the formation of insoluble calcium trioxocarbonate(IV).

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

When CO₂ is continuously bubbled through the solution, the milkiness disappears because soluble calcium hydrogen trioxocarbonate(IV) is formed:

CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)

When the clear solution is heated, it becomes milky again because the soluble calcium hydrogen trioxocarbonate(IV) decomposes:

Ca(HCO₃)₂(aq) → CaCO₃(s) + H₂O(l) + CO₂(g)

Uses of CO₂

1. CO₂ is used in fire extinguishers because it does not support combustion and is denser than air.

2. It is used to produce carbonated (aerated) drinks, giving them a refreshing taste.

3. It is used in the manufacture of compounds such as washing soda (Na₂CO₃) and urea.

4. Solid CO₂ (dry ice) is used for refrigerating goods, such as ice cream.

5. Gaseous CO₂ is used to preserve fruits.

6. It is used as a coolant in nuclear reactors.

7. Yeast and baking powder are used in baking to produce CO₂, which causes dough to rise and makes the loaf light.

                                                   Carbon (II) oxide, CO

Carbon(II) oxide gas is produced during the incomplete combustion of carbon compounds, such as methane, coal and petrol.

For example:

2C₈H₁₈(l) + 17O₂(g) → 16CO(g) + 18H₂O(l)

Carbon(II) oxide is a poisonous gas. It has no colour or odour, making its presence difficult to detect and therefore very dangerous.

Preparation of CO

1. From Carbon(IV) Oxide

Carbon(II) oxide is produced when carbon(IV) oxide is passed over red-hot carbon:

CO₂(g) + C(s) → 2CO(g)

The gas should be prepared in a fume cupboard because carbon(II) oxide is poisonous.

2. Dehydration of Methanoic Acid

Carbon(II) oxide can be prepared by dehydrating methanoic (formic) acid using concentrated tetraoxosulphate(VI) acid as a dehydrating agent:

HCOOH(l) → CO(g) + H₂O(l)

Preparation from Ethanedioic (Oxalic) Acid

Ethanedioic (oxalic) acid can be dehydrated using concentrated H₂SO₄:

H₂C₂O₄ → CO₂ + CO + H₂O

Equal volumes of CO₂ and CO are produced. The gases are passed through concentrated NaOH solution to remove the CO₂, leaving CO.

Physical Properties of CO

1. Carbon(II) oxide is a colourless, odourless and tasteless gas.

2. It is slightly less dense than air.

3. It is slightly soluble in water.

4. It is neutral to litmus.

5. It is a poisonous gas and combines with haemoglobin in the blood to form a stable compound called carboxyhaemoglobin.

6. It is difficult to detect because it has no colour or odour.

Chemical Properties of CO

1. As a Reducing Agent

CO is a strong reducing agent.

For example:

PbO(s) + CO(g) → Pb(s) + CO₂(g)

Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g)

2. Combustion Reaction

CO burns in oxygen with a faint blue flame to form CO₂:

2CO(g) + O₂(g) → 2CO₂(g)

3. Reaction with Hydrogen

CO reacts with hydrogen to form methanol:

CO(g) + 2H₂(g) → CH₃OH(g)

Test for CO

1. Bubble some of the unknown gas through limewater. If the limewater turns milky, the gas is CO₂, not CO.

2. Pass the gas through a flame. CO burns with a blue flame, producing CO₂.

3. The resulting gas can then be passed through limewater; the limewater turns milky.

Uses of CO

1. CO is used as a reducing agent in the extraction of metals from their ores, particularly in metallurgical processes.

2. It is used as a component of certain gaseous fuels.

3. It is used in the manufacture of methanol and other industrial chemicals.

4. It is useful in industrial processes because of its ability to reduce metal oxides.

                                                      FUEL GASES

Fuel gases are gases that can be burnt to produce heat. The main fuel gases are:

1. Producer gas

2. Water gas

                                                 Producer Gas

Producer gas is a mixture of nitrogen (N₂) and carbon(II) oxide (CO).

It is prepared by passing a stream of air through red-hot coke.

2C(s) + O₂(g) → 2CO(g) + heat

The nitrogen in the air does not take part in the reaction and therefore remains mixed with the carbon(II) oxide.

Producer gas contains approximately 67% nitrogen and 33% carbon(II) oxide.

Uses of Producer Gas

1. It is used to heat furnaces, ovens and other industrial equipment.

2. It can be used as a source of nitrogen for the production of ammonia by the Haber process.

                                                   Water Gas

Water gas is a mixture of equal volumes of hydrogen (H₂) and carbon(II) oxide (CO).

It is prepared by passing steam over white-hot coke at about 1,000°C.

C(s) + H₂O(g) → CO(g) + H₂(g)

Both hydrogen and carbon(II) oxide in water gas burn in air and release a large amount of heat.

Uses of Water Gas

1. It is used as an industrial fuel.

2. It can be used in the manufacture of other useful chemicals.

Note: Carbon(II) oxide is poisonous, so water gas must be handled with care.

                                      TRIOXOCARBONATES(IV)

Trioxocarbonate(IV) acid is an acidic compound. It forms two series of salts:

1. Normal trioxocarbonates(IV) — e.g. K₂CO₃

2. Acid trioxocarbonates(IV) — e.g. KHCO₃

Preparation of Soluble Trioxocarbonates(IV)

Sodium, potassium and ammonium trioxocarbonates(IV) are the principal soluble salts.

For example, potassium trioxocarbonate(IV) can be prepared by passing carbon dioxide into potassium hydroxide solution:

2KOH(aq) + CO₂(g) → K₂CO₃(aq) + H₂O(l)

Soluble trioxocarbonates(IV) can also be prepared by heating the corresponding hydrogen trioxocarbonate(IV):

2KHCO₃(s) → K₂CO₃(s) + H₂O(l) + CO₂(g)

Preparation of Insoluble Trioxocarbonates(IV)

Insoluble metallic trioxocarbonates(IV) can be prepared as precipitates by adding a solution of a soluble trioxocarbonate(IV) to a solution of the corresponding metal salt.

CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)

CaCl₂(aq) + 2NaHCO₃(aq) → CaCO₃(s) + 2NaCl(aq) + H₂O(l) + CO₂(g)

                                    HYDROGEN TRIOXOCARBONATE(IV)

Hydrogen trioxocarbonate(IV), HCO₃⁻, is the acid salt of trioxocarbonate(IV) acid.

They are formed when a metal or ammonium radical replaces one of the hydrogen atoms in trioxocarbonate(IV) acid.

They are generally soluble in water and can be prepared by passing carbon dioxide through a cold solution of the corresponding hydroxide or trioxocarbonate(IV).