Metals and Their Compounds

Metals are elements that ionize by losing electrons. They possess unique characteristics that differentiate them from non-metals. Their physical properties depend on the arrangement of atoms or molecules in a crystal lattice, and the type of bonding present in different states of matter. Chemically, the behavior of metals is largely determined by the number of valence electrons in their atoms. Over 75% of known elements are classified as metals.

Physical Properties of Metals

Exceptions

Metals exist as solids at room temperature (except mercury) and form a crystal lattice structure held by strong metallic bonds. Non-metals, in contrast, typically form covalent molecules held by weak intermolecular forces, with exceptions like diamond and graphite (forms of carbon).

Chemical Properties of Metals

Principles of Metal Extraction

Metals are usually found in the Earth's crust as compounds mixed with other substances. To make them useful, they must be extracted. A metal ore is a rock that contains a sufficient concentration of a metal or its compound for extraction to be economically viable.

The extraction method depends on the metal's reactivity:

The Alkali Metals (Group 1 Elements)

General Properties

Sodium

Occurrence

Ores of Sodium

Extraction of Sodium

Sodium was first isolated by Sir Humphrey Davy through electrolysis of fused caustic soda. Two extraction methods include:

Down’s Process

Molten sodium chloride is electrolyzed using a graphite anode and an iron ring cathode. A wire gauge separates the electrodes to avoid sodium and chlorine mixing. A NaCl-KCl-KF mixture is used to lower melting point from 1085 K to ~850-875 K.

Reasons for using a mixture:

Chemical Reactions:

At the cathode:

Na+ + e- → Na

At the anode:

2Cl- → Cl2 + 2e-

Advantages:

Physical Properties of Lithium and Sodium

Chemical Properties of Lithium and Sodium

1. Action of Air

2. Reaction with Oxygen

3. Reaction with Water

Both react with cold water, releasing hydrogen gas.

2Li + 2H2O → 2LiOH + H2
2Na + 2H2O → 2NaOH + H2

4. Reaction with Non-metals

5. Reaction with Ammonia

6. Reducing Action

They reduce metallic oxides/halides to metals (e.g., beryllium, uranium).

7. Reaction with Acids

Violently react to produce hydrogen gas.

8. Reaction with Mercury

Form sodium amalgams like NaHg, Na2Hg, Na3Hg.

9. Solubility in Liquid Ammonia

Dissolve to give a deep blue, conducting solution due to ammoniated electrons.

Uses of Sodium

Compounds of Sodium

Sodium Hydroxide (NaOH)

Physical Properties

Chemical Properties

Uses

Sodium Chloride (NaCl)

Uses

Sodium Tetraoxosulphate (IV) – Na2SO4

Uses

Sodium Trioxonitrate (V) – NaNO3

Uses

Sodium Trioxocarbonate (IV) – Na2CO3

Uses

Alkaline Earth Metals (Group II Elements)

Alkaline earth metals do not exist freely in nature. They are typically found as trioxocarbonates(IV) and tetraoxosulphates(VI). Elements in this group include calcium, magnesium, beryllium, strontium, barium, and radium.

Calcium and magnesium compounds form from rock weathering and are found in soil and fresh water as hydrogen trioxocarbonates(IV), tetraoxosulphates(VI), and chlorides — contributing to water hardness. Bones and teeth contain calcium tetraoxophosphate(V). Beryllium is found in small quantities as tetraoxosilicate minerals and is hard to extract. Strontium and radium are less abundant; radium is radioactive.

Physical Properties of Alkaline Earth Metals

Chemical Properties

Alkaline earth metals tarnish in air, forming oxides. Beryllium forms a protective oxide layer, unlike the others. When heated in air, they burn to produce oxides and nitrides.

2Mg(s) + O2(g) → 2MgO(s)  
3Mg(s) + N2(g) → Mg3N2(s)

Calcium

Occurrence

Calcium is too reactive to occur naturally as a free element. It exists in:

Extraction

Metallic calcium is extracted by electrolysis of fused calcium chloride, a by-product of the Solvay process. Calcium fluoride is added to lower the melting point from 850°C to 650°C. Electrolysis is carried out using a graphite-lined crucible (anode) and an iron rod cathode. Calcium deposits on the cathode; chlorine gas is released at the anode.

Chemical Reactions:

At cathode: Ca2+(aq) + 2e- → Ca(s)  
At anode: 2Cl-(aq) → Cl2(g) + 2e-

Physical Properties of Calcium

Chemical Properties of Calcium

Reacts readily with air to form calcium oxide. In the presence of moisture and CO2, it forms calcium hydroxide and eventually calcium trioxocarbonate(IV):

2Ca(s) + O2(g) → 2CaO(s)  
CaO(s) + H2O(l) → Ca(OH)2(s)  
Ca(OH)2(s) + CO2(g) → CaCO3(s) + H2O(l)

It also reacts with non-metals:

3Ca(s) + N2(g) → Ca3N2(s)  
Ca(s) + Cl2(g) → CaCl2(s)  
Ca(s) + H2(g) → CaH2(s)

With water:

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

With ammonia:

3Ca(s) + 2NH3(g) → Ca3N2(s) + 3H2(g)

With dilute acids:

Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g)

Uses of Calcium

Calcium Oxide (CaO) – Quicklime

Preparation

Produced by heating calcium carbonate above 825°C:

CaCO3(s) → CaO(s) + CO2(g)

Uses

Calcium Hydroxide [Ca(OH)2] – Slaked Lime

Properties

Uses

Calcium Trioxocarbonate(IV) – CaCO3

Occurrence & Preparation

Occurs naturally as limestone, chalk, coral, and in bones and shells. Prepared by reacting CO2 with calcium hydroxide:

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

Properties

Uses

Calcium Chloride – CaCl2

Properties

Uses

Calcium Tetraoxosulphate(VI) – CaSO4

Properties

Uses

Aluminium

Extraction of Aluminium from Bauxite

Aluminium is highly reactive and cannot be extracted from its ore using carbon reduction because the required temperature would be uneconomical. Instead, it is extracted by electrolysis.

Ore

The primary ore of aluminium is bauxite, an impure form of aluminium oxide. Major impurities include iron oxides, silicon dioxide, and titanium dioxide.

Purification of Aluminium Oxide – Bayer Process

Reaction with Sodium Hydroxide

Crushed bauxite is treated with concentrated sodium hydroxide solution under high temperature (140°C–240°C) and pressure (up to 35 atm). Aluminium oxide dissolves to form sodium tetrahydroxoaluminate while impurities remain as solids and are removed by filtration. The residue is called red mud.

Precipitation of Aluminium Hydroxide

The solution is cooled and seeded with aluminium hydroxide crystals, triggering the precipitation of more aluminium hydroxide.

Formation of Pure Aluminium Oxide

Aluminium hydroxide is heated to about 1100–1200°C to form pure aluminium oxide (Al2O3).

Electrolysis of Aluminium Oxide

The purified aluminium oxide is dissolved in molten cryolite (Na3AlF6) and electrolysed. The molten aluminium collects at the bottom as the cathode product, while oxygen is formed at the carbon anodes.

Electrode Reactions

Cathode:

Al3+ + 3e- → Al(l)

Anode:

2O2- → O2(g) + 4e-

Oxygen reacts with the carbon anodes to form carbon dioxide and carbon monoxide:

C(s) + O2(g) → CO2(g)

2C(s) + O2(g) → 2CO(g)

Properties of Aluminium

1. Reaction with Air

When exposed to air, aluminium forms a thin oxide layer that protects it from further corrosion.

At about 800°C:

4Al(s) + 3O2(g) → 2Al2O3(s)

2Al(s) + N2(g) → 2AlN(s)

2. Reaction with Acids

Uses of Aluminium

Common Applications

Tin (Sn)

Occurrence

Tin occurs naturally as Cassiterite or tinstone, with the chemical formula SnO2. Major deposits are found in Malaysia, Bolivia, and Indonesia. In Nigeria, it is found in smaller quantities in graphite rocks and alluvial deposits, especially in Jos.

Extraction of Tin

  1. The ore is first crushed and washed using a strong stream of water.
  2. It is then roasted in air to eliminate impurities like sulphur, arsenic, and antimony as their volatile oxides.
  3. The resulting tin(IV) oxide (SnO2) is reduced by heating it with coke in a reverberatory furnace.

Chemical Equation:

SnO2(s) + 2C(s) → Sn(l) + 2CO(g)

The molten tin is purified further by heating gently on a sloped surface. As it flows down, impurities are oxidized in air and left behind as scum.

Uses of Tin

Iron

Rusting of Iron

Rusting is the corrosion of iron caused by the combined action of oxygen and moisture from the environment. This process is accelerated by pollutants such as carbon dioxide and other gases present in the air. When iron reacts with oxygen and water, it forms a soft, flaky brown substance known as hydrated iron(III) oxide (Fe2O3.xH2O), which breaks off easily.

Reaction: 4Fe(s) + 3O2(g) + 2xH2O(l) → 2Fe2O3.xH2O(s)

Extraction of Iron

Iron is extracted in a blast furnace by reducing its oxide. The iron ore is first roasted in air to produce iron(III) oxide. It is then mixed with coke and limestone and introduced into the furnace. A hot blast of air is blown from the base, initiating the reduction process.

Experiment to Show Air and Moisture Are Required for Rusting

Materials:

Procedure:

  1. Place equal numbers of iron nails in each test tube.
  2. In Tube A, add calcium chloride and seal with a cotton pad and stopper. This removes moisture.
  3. In Tube B, add boiled water and seal with a stopper and rubber tubing to remove air.
  4. In Tube C, add tap water and leave it open to air.

Result:

Only the nails in Tube C rusted, showing that both air and water are necessary for rusting.

Prevention of Rusting

Rusting vs. Burning

Types of Iron

Physical Properties of Iron

Chemical Properties of Iron

Test for Iron(II) Ions

  1. Add NaOH solution: dirty-green precipitate forms, insoluble in excess.
  2. Add NH3: dirty-green precipitate forms, insoluble in excess.
  3. Add potassium hexacyanoferrate(III): forms dark blue precipitate.

Test for Iron(III) Ions

  1. Add NaOH solution: reddish-brown precipitate forms, insoluble in excess.
  2. Add NH3: reddish-brown precipitate forms, insoluble in excess.
  3. Add potassium hexacyanoferrate(II): forms a blue precipitate.

The Transition Metals

Transition elements are found between Group II and Group III of the periodic table and share similar chemical behavior. They are also known as the d-block elements because electrons are progressively added to their d-subshells.

Examples include: Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn.

They are all metals and exhibit typical metallic properties such as high tensile strength and high melting points. Unlike Group I and II metals, transition metals are less reactive and have distinctive characteristics.

General Properties of First Row Transition Elements