Describe the preparation and properties of various compounds of metals (oxides, hydroxides, carbonates, hydrogen carbonates, nitrates, chlorides and sulphates of metals)

Preparation of Oxides

Oxides are prepared by:

1. Direct combination – Metals burn in oxygen to form oxides: $4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O}$ $2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$

2. Heating carbonates – Carbonates decompose when heated to form oxides: $\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$

3. Reacting metals with acids – Oxides can be obtained by reacting metals with acids, then drying the resulting solution.

Properties of Oxides

• Solubility in water: Sodium and potassium oxides react with water to form alkaline solutions (hydroxides): $\text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH}$ $\text{K}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{KOH}$

• Reaction with acids: Metal oxides react with acids to form salt and water: $\text{CuO} + \text{H}_2\text{SO}_4 \rightarrow \text{CuSO}_4 + \text{H}_2\text{O}$ $\text{Fe}_2\text{O}_3 + 6\text{HCl} \rightarrow 2\text{FeCl}_3 + 3\text{H}_2\text{O}$

• Thermal stability: Some metal oxides decompose on heating. For example, silver oxide decomposes: $2\text{Ag}_2\text{O} \rightarrow 4\text{Ag} + \text{O}_2$

Preparation of Hydroxides

1. Soluble hydroxides (alkalis) – Dissolve in water. Sodium hydroxide (caustic soda) and potassium hydroxide (caustic potash) are prepared by electrolysis of salt solutions.

2. Insoluble hydroxides – Prepared by mixing solutions of a soluble salt with a soluble hydroxide (precipitation method): $\text{Cu(NO}_3)_2(\text{aq}) + \text{Ca(OH)}_2(\text{aq}) \rightarrow \text{Cu(OH)}_2(\text{s}) + \text{Ca(NO}_3)_2(\text{aq})$

Properties of Hydroxides

• Solubility: Sodium, potassium, and calcium hydroxides dissolve in water. Most other metal hydroxides are insoluble.

• Bitter taste and slippery feel: All hydroxides have a bitter taste and feel slippery.

• Neutralization: Hydroxides neutralize acids to form salt and water: $\text{NaOH} + \text{HCl} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

• Indicator tests: Hydroxides turn red litmus blue and phenolphthalein pink.

Preparation of Carbonates

1. Sodium and potassium carbonates – Prepared by the Solvay process:

   • Carbon dioxide is passed through a solution of ammonia and brine
   • Sodium hydrogen carbonate precipitates out
   • Heating produces sodium carbonate: $2\text{NaHCO}_3 \rightarrow \text{Na}_2\text{CO}_3 + \text{CO}_2 + \text{H}_2\text{O}$

2. Other carbonates – Can be prepared by reacting metal oxides, hydroxides, or metals with carbon dioxide: $\text{CaO} + \text{CO}_2 \rightarrow \text{CaCO}_3$

Preparation of Hydrogen Carbonates

Hydrogen carbonates (bicarbonates) are formed when carbon dioxide is passed through solutions of hydroxides or carbonates: $\text{NaOH} + \text{CO}_2 \rightarrow \text{NaHCO}_3$ $\text{Na}_2\text{CO}_3 + \text{CO}_2 + \text{H}_2\text{O} \rightarrow 2\text{NaHCO}_3$

Properties of Carbonates and Hydrogen Carbonates

• Solubility: All carbonates are insoluble in water except those of sodium, potassium, and ammonium.

• Reaction with acids: All carbonates and hydrogen carbonates react with dilute acids to produce carbon dioxide, water, and a salt: $\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{CO}_2 + \text{H}_2\text{O}$ $\text{NaHCO}_3 + \text{HCl} \rightarrow \text{NaCl} + \text{CO}_2 + \text{H}_2\text{O}$

• Action of heat: Carbonates of sodium and potassium do not decompose on heating. Other carbonates decompose to form oxide and carbon dioxide: $\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$

Test for Carbonates and Hydrogen Carbonates

• Add dilute acid – effervescence of carbon dioxide gas confirms carbonate or hydrogen carbonate.
• Add magnesium sulphate solution – white precipitate indicates carbonate (no precipitate for hydrogen carbonate).

Preparation of Nitrates

1. Neutralization – React nitric acid with metals, oxides, hydroxides, or carbonates: $\text{NaOH} + \text{HNO}_3 \rightarrow \text{NaNO}_3 + \text{H}_2\text{O}$

2. Action of nitric acid on metals – Only magnesium, zinc, lead, and copper react with dilute nitric acid: $\text{Mg} + 2\text{HNO}_3 \rightarrow \text{Mg(NO}_3)_2 + \text{H}_2$

Properties of Nitrates

• Solubility: All nitrates are soluble in water.

• Action of heat: Different nitrates decompose differently on heating:

  • Sodium and potassium nitrates give nitrites: $2\text{KNO}_3 \rightarrow 2\text{KNO}_2 + \text{O}_2$

  • Most metal nitrates give oxides: $2\text{Pb(NO}_3)_2 \rightarrow 2\text{PbO} + 4\text{NO}_2 + \text{O}_2$

  • Silver and mercury nitrates give metals: $2\text{AgNO}_3 \rightarrow 2\text{Ag} + 2\text{NO}_2 + \text{O}_2$

• Brown ring test: Add freshly prepared ferrous sulphate solution, then carefully add concentrated sulphuric acid. A brown ring at the junction confirms nitrate.

Preparation of Chlorides

1. Soluble chlorides – React metals, metal oxides, hydroxides, or carbonates with dilute hydrochloric acid: $\text{ZnO} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2\text{O}$ $\text{Mg} + 2\text{HCl} \rightarrow \text{MgCl}_2 + \text{H}_2$ $\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{CO}_2 + \text{H}_2\text{O}$

2. Insoluble chlorides – Prepared by precipitation: $\text{Pb(NO}_3)_2(\text{aq}) + 2\text{NaCl}(\text{aq}) \rightarrow \text{PbCl}_2(\text{s}) + 2\text{NaNO}_3(\text{aq})$ $\text{AgNO}_3(\text{aq}) + \text{NaCl}(\text{aq}) \rightarrow \text{AgCl}(\text{s}) + \text{NaNO}_3(\text{aq})$

Properties of Chlorides

• Solubility: Most chlorides are soluble. Silver chloride, lead(II) chloride, and mercury(I) chloride are insoluble.

• Test for chlorides: Add silver nitrate solution to a solution of chloride – white precipitate of silver chloride forms. The precipitate dissolves in dilute ammonia solution.

Preparation of Sulphates

1. React metals, oxides, hydroxides, or carbonates with dilute sulphuric acid: $\text{ZnO} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2\text{O}$ $\text{Mg} + \text{H}_2\text{SO}_4 \rightarrow \text{MgSO}_4 + \text{H}_2$

2. Direct combination (for some metals): $2\text{Fe} + 3\text{O}_2 + 2\text{H}_2\text{SO}_4 \rightarrow \text{Fe}_2(\text{SO}_4)_3 + 2\text{H}_2\text{O}$

Properties of Sulphates

• Solubility: Most sulphates are soluble. Barium sulphate, lead(II) sulphate, and calcium sulphate are insoluble.

• Reaction with acids: Sulphates do not generally react with acids.

• Test for sulphates: Add barium chloride solution – white precipitate of barium sulphate forms: $\text{BaCl}_2 + \text{Na}_2\text{SO}_4 \rightarrow \text{BaSO}_4(\text{s}) + 2\text{NaCl}$

• Hydrated crystals: Many sulphates form crystals with water of crystallization (e.g., copper(II) sulphate pentahydrate, $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$).

Question: Describe how you would prepare pure lead(II) chloride in the laboratory.

Answer:

1. Method: Use precipitation because lead(II) chloride is insoluble.

2. Procedure:

   • Dissolve lead(II) nitrate in water in a beaker
   • Dissolve sodium chloride in water in another beaker
   • Mix the two solutions together
   • A white precipitate of lead(II) chloride forms: $\text{Pb(NO}_3)_2(\text{aq}) + 2\text{NaCl}(\text{aq}) \rightarrow \text{PbCl}_2(\text{s}) + 2\text{NaNO}_3(\text{aq})$

3. Separation: Filter the mixture to collect the precipitate

   • Wash with distilled water to remove impurities
   • Dry on filter paper

4. Confirmation: The white precipitate is insoluble in cold water but dissolves in hot water. When cooled, it crystallizes out again.

In Tanzania, the knowledge of metal compounds is applied in water treatment. Lime (calcium hydroxide) is used to soften hard water by removing calcium and magnesium carbonates. Copper sulphate is used in fish ponds to control algae growth and treat fungal infections in fish. Additionally, sodium carbonate (washing soda) is commonly used in laundry to remove grease and soften water—families in Tanzania purchase it in markets for cleaning clothes.