Chemical Reactions and Equations

1. What Is a Chemical Reaction?

Chemical Reaction: A process in which one or more substances undergo a chemical change to form new substances with different properties and identities.

During a chemical reaction, the original substances are converted into new substances. The products formed have properties different from those of the reactants.

Common Indicators of a Chemical Reaction

Change in state.
Change in colour.
Evolution of a gas.
Change in temperature.
Formation of a precipitate.

Observation	Example
Change in colour	Green ferrous sulphate crystals change colour on heating.
Evolution of gas	Zinc reacts with dilute sulphuric acid and releases hydrogen gas.
Change in temperature	Zinc reacting with acid or calcium oxide reacting with water releases heat.
Formation of solid	Some reactions produce an insoluble solid called a precipitate.

Magnesium Ribbon Experiment

Magnesium ribbon is cleaned with sandpaper before burning because it develops a thin protective layer of magnesium oxide on its surface.

1. Clean the magnesium ribbon using sandpaper.

2. Hold it with tongs and burn it in air.

3. It burns with a dazzling white flame.

4. A white powder of magnesium oxide is formed.

2Mg(s) + O₂(g) → 2MgO(s)

Magnesium oxide is a new substance with properties different from magnesium and oxygen.

2. Chemical Equations

A chemical reaction can be represented using chemical symbols and formulae. This representation is called a chemical equation.

Skeletal Equation

A skeletal equation is an unbalanced chemical equation in which the number of atoms of each element may not be equal on both sides.

Mg + O₂ → MgO

Balanced Chemical Equation

A balanced chemical equation contains an equal number of atoms of every element on both the reactant and product sides.

2Mg + O₂ → 2MgO

Law of Conservation of Mass: Mass can neither be created nor destroyed during a chemical reaction. Therefore, the total mass of reactants equals the total mass of products.

Rules for Balancing Equations

1. Write the correct skeletal equation.

2. Count the atoms of every element on both sides.

3. Start balancing the compound containing the maximum number of atoms.

4. Adjust only the coefficients placed before formulae.

5. Never change the subscripts inside a chemical formula.

6. Check that all atoms are equal on both sides.

Changing a subscript changes the substance itself. Only coefficients should be changed while balancing an equation.

3. Information Conveyed by a Chemical Equation

Physical State Symbols

Symbol	Meaning
(s)	Solid
(l)	Liquid
(g)	Gas
(aq)	Aqueous solution dissolved in water

Reaction Conditions

Conditions such as temperature, pressure, sunlight, electricity or catalysts may be written above or below the reaction arrow.

Δ: Heating
Sunlight: Photochemical reaction
Electricity: Electrolytic decomposition
Chlorophyll/Sunlight: Used in photosynthesis
Pressure: May be written when required

Water vapour is represented as H₂O(g), while liquid water is represented as H₂O(l).

4. Combination Reactions

Combination Reaction: A reaction in which two or more reactants combine to form a single product.

A + B → AB

Quick Lime and Water

Calcium oxide, commonly called quick lime, reacts vigorously with water to form calcium hydroxide or slaked lime.

CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

Since heat is released, this is an exothermic combination reaction.

Whitewashing Reaction

Slaked lime is applied to walls during whitewashing. It reacts slowly with carbon dioxide in the atmosphere to form a hard, shiny layer of calcium carbonate.

Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

Calcium carbonate is also the main chemical component of marble. The whitewashed surface becomes shiny after approximately two to three days.

5. Exothermic and Endothermic Reactions

Exothermic Reaction

A reaction in which heat is released into the surroundings.

Releases energy.
Surroundings become warmer.
Often produces a rise in temperature.

Endothermic Reaction

A reaction in which energy is absorbed from the surroundings.

Absorbs heat, light or electrical energy.
Surroundings may become cooler.
Often requires continuous energy input.

Examples of Exothermic Reactions

Respiration.
Burning of natural gas.
Reaction of calcium oxide with water.
Decomposition of vegetable matter into compost.

Respiration as an Exothermic Process

C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy

Energy is released when glucose reacts with oxygen inside body cells.

6. Decomposition Reactions

Decomposition Reaction: A reaction in which a single reactant breaks down into two or more simpler products.

AB → A + B

Types of Decomposition Reactions

Type	Energy Source	Example
Thermal Decomposition	Heat	Decomposition of ferrous sulphate or limestone
Electrolytic Decomposition	Electricity	Electrolysis of water
Photolytic Decomposition	Sunlight	Decomposition of silver chloride

Thermal Decomposition of Ferrous Sulphate

Ferrous sulphate crystals are green. On heating, they first lose water and then decompose to form ferric oxide and sulphur oxides.

2FeSO₄(s) → Fe₂O₃(s) + SO₂(g) + SO₃(g)

Green crystals change colour.
A brown solid of ferric oxide is formed.
Gases with the smell of burning sulphur are released.

Thermal Decomposition of Lead Nitrate

2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

Brown fumes of nitrogen dioxide are observed.

Thermal Decomposition of Limestone

CaCO₃(s) → CaO(s) + CO₂(g)

This reaction is important in the manufacture of cement.

Electrolysis of Water

2H₂O(l) → 2H₂(g) + O₂(g)

Hydrogen and oxygen are produced in the volume ratio of 2:1 because water contains two hydrogen atoms for every one oxygen atom.

Photolytic Decomposition

2AgCl(s) → 2Ag(s) + Cl₂(g)

2AgBr(s) → 2Ag(s) + Br₂(g)

Silver chloride and silver bromide decompose in sunlight to form grey silver. These reactions are used in traditional black-and-white photography.

7. Displacement Reactions

Displacement Reaction: A reaction in which a more reactive element displaces a less reactive element from its compound.

A + BC → AC + B

Iron Nail and Copper Sulphate Experiment

When an iron nail is placed in blue copper sulphate solution, iron displaces copper because iron is more reactive than copper.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Observations

A reddish-brown coating of copper forms on the iron nail.
The blue colour of copper sulphate solution fades.
The solution becomes pale green due to iron sulphate formation.

Zinc and lead can also displace copper from copper salts because they are more reactive than copper.

8. Double Displacement and Precipitation Reactions

Double Displacement Reaction: A reaction in which two compounds exchange ions to form two new compounds.

AB + CD → AD + CB

Precipitation Reaction

A precipitation reaction is a reaction in which an insoluble solid, called a precipitate, is formed.

Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)

Barium sulphate is formed as a white precipitate.
Sodium chloride remains dissolved in the solution.
The reaction involves exchange of ions.

The insoluble solid formed during a chemical reaction is called a precipitate.

9. Oxidation, Reduction and Redox Reactions

Oxidation

Oxidation is the gain of oxygen or loss of hydrogen by a substance.

Reduction

Reduction is the loss of oxygen or gain of hydrogen by a substance.

Redox Reaction

A reaction in which oxidation and reduction occur simultaneously is called a redox reaction.

CuO + H₂ → Cu + H₂O

Substance	Change	Process
CuO	Loses oxygen and becomes copper	Reduction
H₂	Gains oxygen and becomes water	Oxidation

Heating Copper Powder

Copper powder reacts with oxygen on heating and develops a black coating of copper(II) oxide.

2Cu + O₂ → 2CuO

The black colour indicates the formation of copper(II) oxide.

10. Effects of Oxidation in Daily Life

Corrosion

Corrosion: The slow deterioration of metals due to reactions with substances in the environment, such as oxygen, moisture and acids.

Metal	Corrosion Product	Appearance
Iron	Rust	Reddish-brown coating
Silver	Silver sulphide or related compounds	Black coating
Copper	Basic copper carbonate	Green coating

Examples of Corrosion

Rusting of iron objects.
Blackening of silver articles.
Formation of a green coating on copper surfaces.

Rancidity

Rancidity: The oxidation of fats and oils in food, causing unpleasant smell and taste.

Rancidity reduces the quality of food and may make it unsuitable for consumption.

Methods to Prevent Rancidity

Adding antioxidants to food.
Keeping food in airtight containers.
Refrigerating food when appropriate.
Packaging food in an inert atmosphere.
Flushing packets of chips with nitrogen gas.

Nitrogen is relatively unreactive and prevents oxygen from coming into contact with fats and oils in packaged food.

11. Summary of Reaction Types

Reaction Type	General Form	Example
Combination	A + B → AB	CaO + H₂O → Ca(OH)₂
Decomposition	AB → A + B	CaCO₃ → CaO + CO₂
Displacement	A + BC → AC + B	Fe + CuSO₄ → FeSO₄ + Cu
Double Displacement	AB + CD → AD + CB	Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl
Redox	Oxidation and reduction together	CuO + H₂ → Cu + H₂O

12. Quick Revision