Chemical Reactions and Equations

Unlocking the mysteries of chemical transformations

What is a Chemical Reaction?

A chemical reaction occurs whenever a chemical change takes place, altering the nature and identity of the initial substance.

Examples from daily life include:

How to Determine if a Chemical Reaction Has Taken Place:

Reactants and Products:

Reactants:

Substances that undergo chemical change in a reaction (written on the Left-Hand Side - LHS).

Products:

New substances formed during the reaction (written on the Right-Hand Side - RHS).

An arrow (→) is placed between reactants and products, pointing towards the products and indicating the direction of the reaction. A plus sign (+) separates multiple reactants or products.

Word Equations vs. Chemical Equations:

Word-equation:

A description of a chemical reaction in sentence form, written in a shorter word-equation format.

Example: Magnesium + Oxygen → Magnesium oxide

Chemical equation:

A more concise and useful representation using chemical formulae instead of words.

Example: Mg + O₂ → MgO

Skeletal vs. Balanced Chemical Equations:

Skeletal chemical equation:

An unbalanced chemical equation where the number of atoms of each element is not the same on both sides of the equation. It's unbalanced because the mass is not the same on both sides.

Example: Mg + O₂ → MgO

Balanced chemical equation:

An equation where the number of atoms of each element is the same on both sides, adhering to the Law of Conservation of Mass. This law states that mass can neither be created nor destroyed in a chemical reaction; thus, the total mass and number of atoms of elements must be equal before and after the reaction.

Steps to Balance a Chemical Equation (Hit-and-Trial Method):

  1. Draw boxes around each formula; do not change anything inside the boxes.
  2. List the number of atoms of different elements on the LHS and RHS of the unbalanced equation.
  3. Start balancing with the compound containing the maximum number of atoms, and within that, the element with the maximum number of atoms. Use coefficients (e.g., '4 H₂O', not 'H₂O₄').
  4. Balance other elements one by one.
  5. Check the correctness by counting atoms of each element on both sides.

Writing Symbols of Physical States and Reaction Conditions:

To make an equation more informative, physical states are mentioned with chemical formulae:

Physical states are usually not included unless necessary.

Reaction conditions (temperature, pressure, catalyst) can be indicated above and/or below the arrow.

Types of Chemical Reactions

Chemical reactions involve the breaking and making of bonds between atoms to produce new substances. Atoms of one element do not change into those of another.

1. Combination Reaction:

A reaction in which two or more substances (elements or compounds) combine to form a single product.

Examples:

  • Formation of slaked lime: CaO(s) + H₂O(l) → Ca(OH)₂(aq)
  • Burning of coal: C(s) + O₂(g) → CO₂(g)
  • Formation of water: 2H₂(g) + O₂(g) → 2H₂O(l)

Often release a large amount of heat, making them exothermic reactions.

Exothermic and Endothermic Reactions:

Exothermic reactions:

Reactions in which heat is released along with the formation of products.

Examples:

  • Burning of natural gas: CH₄ + 2O₂ → CO₂ + 2H₂O
  • Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy
  • Decomposition of vegetable matter into compost.

Endothermic reactions:

Reactions in which energy is absorbed (in the form of heat, light, or electricity) for breaking down reactants.

2. Decomposition Reaction:

A reaction where a single reactant breaks down to give two or more simpler products. This is the opposite of a combination reaction.

Types based on energy supplied:

Thermal decomposition (by heating):

  • Ferrous sulphate decomposition: 2FeSO₄(s) --(Heat)--> Fe₂O₃(s) + SO₂(g) + SO₃(g)
  • Limestone (calcium carbonate) decomposition: CaCO₃(s) --(Heat)--> CaO(s) + CO₂(g)
  • Lead nitrate decomposition: 2Pb(NO₃)₂(s) --(Heat)--> 2PbO(s) + 4NO₂(g) + O₂(g)

Decomposition by light:

  • Silver chloride: 2AgCl(s) --(Sunlight)--> 2Ag(s) + Cl₂(g)
  • Silver bromide: 2AgBr(s) --(Sunlight)--> 2Ag(s) + Br₂(g) (Used in black and white photography).

Decomposition by electricity (Electrolysis):

  • Water: 2H₂O(l) --(Electricity)--> 2H₂(g) + O₂(g)

3. Displacement Reaction:

A reaction in which one element displaces or removes another element from its compound.

Principle: A more reactive element displaces a less reactive element.

Examples:

  • Iron displacing copper from copper sulphate: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
  • Zinc displacing copper from copper sulphate: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
  • Lead displacing copper from copper chloride: Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

4. Double Displacement Reaction:

Reactions in which there is an exchange of ions between the reactants.

Often produce a precipitate, an insoluble substance formed during the reaction. Reactions producing a precipitate are called precipitation reactions.

Example:

  • Sodium sulphate and barium chloride: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq). The white precipitate is barium sulphate (BaSO₄).

Oxidation and Reduction (Redox Reactions):

Oxidation:

  • If a substance gains oxygen during a reaction, it is said to be oxidised.
  • If a substance loses hydrogen during a reaction, it is oxidised.

Reduction:

  • If a substance loses oxygen during a reaction, it is said to be reduced.
  • If a substance gains hydrogen during a reaction, it is reduced.

Redox reactions:

Reactions where one reactant gets oxidised while the other gets reduced simultaneously.

Examples:

  • CuO + H₂ --(Heat)--> Cu + H₂O: Copper(II) oxide (CuO) is reduced (loses oxygen), and hydrogen (H₂) is oxidised (gains oxygen).
  • ZnO + C → Zn + CO: Carbon (C) is oxidised to CO, and zinc oxide (ZnO) is reduced to Zn.
  • MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂: HCl is oxidised to Cl₂, and MnO₂ is reduced to MnCl₂.

Effects of Oxidation Reactions in Everyday Life

1. Corrosion:

Definition: The process when a metal is attacked by substances around it such as moisture, acids, etc.. This causes the metal to corrode.

Examples:

  • Rusting of iron: Iron articles get coated with a reddish-brown powder when exposed to humid atmosphere. This is a serious problem causing damage to car bodies, bridges, iron railings, and ships.
  • Black coating on silver.
  • Green coating on copper.

2. Rancidity:

Definition: When fats and oils are oxidised, they become rancid, and their smell and taste change.

Prevention:

  • Adding antioxidants (substances that prevent oxidation) to foods containing fats and oils.
  • Keeping food in air-tight containers to slow down oxidation.
  • Flushing bags of chips with nitrogen gas by manufacturers to prevent oxidation.