Chemical Reactions

A chemical reaction rearranges atoms from reactants into new products, breaking old bonds and forming new ones. Mass is conserved (atoms are neither created nor destroyed), so every equation must be balanced. The eight fundamental reaction types below cover the vast majority of chemistry encountered from high school through undergraduate courses.

Synthesis (Combination) A + B → AB

In a synthesis reaction two or more substances combine to form a single, more complex product. The reactants can be elements or compounds.

Problem Set

Decomposition AB → A + B

A single compound breaks down into two or more simpler substances. Energy input (heat, light, or electricity) is usually required.

Problem Set

Single Displacement A + BC → AC + B

A more reactive free element displaces a less reactive element from a compound. Reactivity (activity) series governs whether the reaction occurs.

Problem Set

Double Displacement (Metathesis) AB + CD → AD + CB

Ions from two compounds exchange partners. The reaction is driven by formation of a precipitate, a gas, or a weakly ionised compound (like water).

Problem Set

Combustion fuel + O₂ → CO₂ + H₂O

Complete combustion occurs with excess oxygen producing CO₂ and H₂O. Incomplete combustion (limited O₂) yields toxic CO and particulate soot (C).

Problem Set

Acid-Base (Neutralisation) acid + base → salt + H₂O

An acid donates H⁺ to a base (Brønsted-Lowry definition). The resulting ionic compound is a salt, and water is usually produced. Net ionic equation: $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$.

Problem Set

Redox (Oxidation-Reduction) electron transfer

Oxidation is loss of electrons (OIL); reduction is gain of electrons (RIG). The species that loses electrons is the reducing agent; the one that gains electrons is the oxidising agent. Oxidation numbers track electron bookkeeping.

Problem Set

Precipitation ions → solid↓

When two aqueous solutions are mixed, ions combine to form an insoluble ionic compound that separates from solution as a solid precipitate. Solubility rules determine which combinations precipitate.

Problem Set

Master Balancing Guide

A balanced equation has equal numbers of each atom on both sides and equal total charge. Use coefficients only — never change subscripts.

Step-by-Step Algorithm

1. Write the unbalanced skeleton — correct formulae for all reactants and products with an arrow.
2. Count atoms of each element on each side. List the imbalances.
3. Balance metals first, then non-metals, then hydrogen, and save oxygen for last.
4. Use the smallest integer coefficients that balance every element simultaneously. Start with the most complex formula and work outwards.
5. Verify — recount all atoms and charge. Reduce coefficients by their GCF if needed.

Worked Examples

Oxidation Number Method (for Redox)

1. Assign oxidation numbers to every atom using standard rules (pure element = 0; O usually −2; H usually +1; sum of compound = 0).
2. Identify which atoms change oxidation number. Calculate the magnitude of change for each.
3. Multiply coefficients so that the total electrons lost (oxidation) equals total electrons gained (reduction) — this is the electron balance.
4. Balance the remaining atoms (not involved in electron transfer) by inspection. Check charge balance in ionic equations.
5. Verify atom counts and charge on both sides.

Example: Balance $\text{KMnO}_4 + \text{HCl} \rightarrow \text{KCl} + \text{MnCl}_2 + \text{Cl}_2 + \text{H}_2\text{O}$

Mn goes +7 → +2 (gain 5e⁻, reduced). Cl goes −1 → 0 (loses 1e⁻ per Cl atom, oxidised). To equalise: 1 Mn × 5e⁻ = 5 Cl × 1e⁻. Coefficient 2 KMnO₄ and 10 HCl for the redox Cl; then balance remaining Cl and H by inspection:

2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 5Cl₂ + 8H₂O

Reaction Energy

Every reaction involves breaking existing bonds (requires energy) and forming new bonds (releases energy). The net difference is the enthalpy change $\Delta H$.

• Exothermic ($\Delta H < 0$): products have lower energy than reactants; heat is released to surroundings. Examples: combustion, neutralisation, rusting.
• Endothermic ($\Delta H > 0$): products have higher energy than reactants; heat is absorbed from surroundings. Examples: photosynthesis, dissolving ammonium nitrate, cooking.
• Activation energy ($E_a$): minimum energy needed to initiate the reaction — the "energy hill" that must be surmounted regardless of whether the reaction is exo- or endothermic.
• Catalysts lower $E_a$ without changing $\Delta H$, speeding up reactions by providing an alternative pathway.

The enthalpy change can be calculated from bond energies: $$\Delta H = \sum E_\text{bonds broken} - \sum E_\text{bonds formed}$$ or from Hess's Law by combining known reaction enthalpies. Standard enthalpy of formation values ($\Delta H_f^\circ$) allow calculation of any reaction enthalpy: $$\Delta H_\text{rxn}^\circ = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})$$