Metals

Metals 🪨

Alloys ⚙️

Definition: An alloy is a mixture of two or more metals or a metal with a non-metal. The combination of alloys is physical. Alloys are harder than the metals from which they are made. Alloying a metal is one way of increasing its strength. 💪

Preparation of Alloys 🔥

The mixture is usually heated under controlled temperature. The molten mixture is then allowed to cool and solidify. ❄️

Advantages of Alloys 🌟

They are flexible in use. 🔄
They usually have improved appearance. ✨
They are durable and reliable. 🛡️
They have increased resistance to corrosion. 🚫💧

Examples of Alloys 🏆

Alloy	Typical Composition	Particular Property	Typical Use
Brass	Copper (Cu) - 70%, Zinc (Zn) - 30%	Harder than pure copper, Resistant to corrosion	Used to make musical instruments
Bronze	Copper (Cu) - 90%, Tin (Sn) - 10%	Harder than pure copper	Used to make trophies, medallions
Mild Steel	Iron (Fe) - 99.7%, Carbon (C) - 0.3%	Stronger and harder than pure iron	Used in bridges, tools, and car bodies
Stainless Steel	Iron (Fe) - 70%, Chromium (Cr) - 20%, Nickel (Ni) - 10%	Does not rust	Used in cutlery, e.g., razor blades, knives
Solder	Tin (Sn) - 50%, Lead (Pb) - 50%	Lower melting point than either tin or lead	Used for soldering and welding

The Reactivity Series ⚡

Alternative Term: Activity Series. Reactivity is a list of metals with the most reactive metal at the top and the least reactive metal at the bottom. 📊

Order of Reactivity 🔝

Potassium 🍌
Sodium 🧂
Calcium 🥛
Magnesium ⚗️
Aluminium 🥫
Zinc ⚙️
Iron 🛠️
Lead 🖊️
Hydrogen 💧
Copper 🪙
Mercury 🌡️
Silver 🥈
Gold 🥇

Reaction of Metals with Water or Steam

Metal	Symbol	Observation / Equation
Potassium	K	Reacts very violently with cold water to produce potassium hydroxide and hydrogen gas. 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Reaction of Metals with Dilute Hydrochloric Acid

All metals above hydrogen in the reactivity series react with dilute hydrochloric acid to form a salt and hydrogen gas. Metals below hydrogen do not react with dilute acids.

Metal	Symbol	Observation / Equation
Potassium	K	Explosive reaction K(s) + HCl(aq) → KCl(aq) + H₂(g)
Sodium	Na	Explosive reaction Na(s) + HCl(aq) → NaCl(aq) + H₂(g)
Calcium	Ca	Reacts vigorously Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g)

Reaction of Metals with Oxygen

Metal	Symbol	Behavior / Equation
Sodium	Na	Catches fire with little heating; burns fiercely with a bright yellow flame 4Na(s) + 2O₂(g) → 2Na₂O(s)
Magnesium	Mg	Catches fire easily; burns with a blinding white flame 2Mg(s) + O₂(g) → 2MgO(s)
Iron	Fe	Glows brightly, giving off yellow sparks 3Fe(s) + 2O₂(g) → Fe₃O₄(s)

Displacement of Metals from their Compounds

A metal high in the reactivity series will always displace a less reactive metal from the solution of its compound.

Example 1: Iron displaces copper from copper (II) sulphate solution:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Example 2: Copper displaces silver from silver nitrate solution, turning it from colorless to blue:

Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s)

Recycling of Metals ♻️

Conserves natural resources, ensuring metal reserves last longer. 🌍
Reduces the need for mining, thus lowering environmental pollution. 🌱
Provides effective waste disposal, reducing landfill space requirements. 🗑️
However, recycling can be costly, especially for cheaper metals. 💰

Extraction of Metals ⛏️

Most metals occur as ores in the Earth's crust. An ore is a compound from which a metal can be extracted. 🌋

Methods of Extraction

Electrolysis: Used for highly reactive metals such as potassium, sodium, and calcium.
Reduction by Coke: Used for moderately reactive metals like zinc and iron.
Thermal Decomposition: Suitable for less reactive metals, such as copper and silver.

Stability of Compounds

Compounds of very reactive metals, such as potassium and sodium, cannot be decomposed by heating or reduction using carbon. Less reactive metal compounds are more easily decomposed and can be reduced by carbon.

Effects of Heat on Carbonates

Group I carbonates (e.g., potassium, sodium) are extremely stable and do not decompose when heated. Group II carbonates (e.g., calcium, magnesium) decompose to oxides and carbon dioxide when heated:

CaCO₃(s) → CaO(s) + CO₂(g)

Extraction of Aluminium 🥡

Aluminium is extracted by electrolysis from bauxite (Al₂O₃) in molten cryolite. This lowers the melting point and reduces energy costs.

Key Reactions

At the Cathode: Al³⁺ + 3e^- → Al
At the Anode: 2O^2- → O₂ + 4e^-

Uses of Aluminium 🥡

Overhead electrical cables: Good conductor of electricity. ⚡
Cooking utensils: Good conductor of heat and resistant to corrosion. 🍳
Aeroplane parts: Lightweight with high strength. ✈️
Food containers and drink cans: Non-toxic and corrosion-resistant. 🥤

Extraction of Copper 🪙

Copper is extracted from copper (I) sulphide by thermal decomposition in air. The impure product, blister copper, is purified by electrolysis.

Uses of Copper 🛠️

Electrical cables: Excellent conductor of electricity. ⚡
Alloys: Used to create bronze and brass. 🥇
Jewelry: Used in ornaments and coins. 💍

Extraction of Zinc ⚙️

Zinc is extracted from zinc blend (ZnS) through roasting and reduction. It is then purified for various uses:

Galvanizing steel to prevent rusting.
Dry cell batteries.

Extraction of Iron 🛠️

Iron is extracted from hematite (Fe₂O₃) in a blast furnace. Carbon monoxide acts as a reducing agent:

Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)

Uses of Slag from Blast Furnace 🏗️

Manufacturing cement. 🏭
Road construction. 🛣️

Effects of Heat on Nitrates

Nitrates of Group I elements, such as potassium and sodium, decompose into nitrites and oxygen when heated. For example:

2KNO₃(s) → 2KNO₂(s) + O₂(g)

Nitrates of Group II elements, such as calcium, magnesium, and transition metals, decompose to metal oxides, nitrogen dioxide, and oxygen:

2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂

Aluminium

Abundance: Aluminium is the most abundant metal in the earth's crust.

Ores of Aluminium

Bauxite (Aluminium oxide - Al₂O₃): A stable compound, it requires electrolysis for extraction.

Extraction Process

Aluminium is extracted by electrolysis of molten bauxite in the presence of cryolite to lower its melting point. The process occurs in a graphite-lined electrolytic cell, with aluminium collecting at the cathode:

At the Cathode: Al³⁺ + 3e^- → Al
At the Anode: 2O^2- → O₂ + 4e^-

The oxygen produced at the anode reacts with the graphite electrodes, forming carbon dioxide, which causes gradual wear on the anodes.

Uses of Aluminium

Overhead electrical cables due to its conductivity and light weight.
Cooking utensils and food wrappers because it is resistant to corrosion and non-toxic.
Airplane and vehicle parts for its strength and low density.
Used in alloys, like duralumin, for additional strength.

Exercise on Aluminium

Example: Explain why aluminium does not corrode easily. Discuss its uses and associated properties.

Extraction of Copper 🪙

Ores of Copper

Malachite (Copper (II) carbonate - CuCO₃)
Copper Pyrites (Copper (I) sulphide - Cu₂S)

Extraction Process

Copper is extracted from copper (I) sulphide through thermal decomposition in air. Impurities such as silver and gold are removed during the electrolysis purification process.

Cu₂S(l) + O₂(g) → 2Cu(s) + SO₂(g)

Purification of Copper

The impure copper anode loses mass as copper ions dissolve into solution, while pure copper is deposited on the cathode.

Uses of Copper 🛠️

Electrical cables due to excellent conductivity.
Alloy production for bronze and brass.
Jewelry and coins, because of its appearance and durability.

Extraction of Zinc ⚙️

Ores of Zinc

Zinc Blende (Zinc sulphide - ZnS)
Calamine (Zinc carbonate - ZnCO₃)

Extraction Process

Zinc blende is roasted in air to form zinc oxide, which is then reduced to zinc metal by carbon monoxide.

2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)

ZnO(s) + CO(g) → Zn(s) + CO₂(g)

Uses of Zinc

Galvanizing steel to prevent rusting.
Alloy production, such as brass (zinc and copper).
Dry cell battery manufacturing.

Extraction of Iron 🛠️

Ores of Iron

Hematite (Iron (III) oxide - Fe₂O₃)
Magnetite (Iron (II, III) oxide - Fe₃O₄)
Siderite (Iron (II) carbonate - FeCO₃)

Extraction Process

Iron is extracted in a blast furnace where carbon monoxide acts as a reducing agent for iron oxide:

Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)

Role of Limestone in Blast Furnace

Limestone decomposes to form calcium oxide, which reacts with impurities to form slag. This slag floats on the molten iron, protecting it from oxidation.

CaCO₃(s) → CaO(s) + CO₂(g)

CaO(s) + SiO₂(s) → CaSiO₃(l)

Uses of Slag

Road construction material.
Used in cement production.

Types of Iron

Cast Iron: Contains impurities and is brittle.
Wrought Iron: The purest form, malleable and ductile.

Exercise Questions ❓

1. Describe the reaction of iron with cold water and steam. Indicate if the reaction takes place. 💧
2. List and describe two methods for preventing rusting of iron. 🛡️
3. Explain the importance of recycling metals. ♻️