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Chemical Equations and Reactions

Subject: Chemistry
Topic: 3
Cambridge Code: 0620 / 0971 / 5070

Chemical Equations​

Chemical equation - Representation of chemical reaction

Word Equation​

Reactants → Products

Example: Hydrogen + Oxygen → Water

Symbol Equation​

Unbalanced: 2H2+O2→H2O2H_2 + O_2 → H_2O

Balanced: 2H2+O2→2H2O2H_2 + O_2 → 2H_2O

Balancing Equations​

Rules:

  • Atoms must be equal on both sides
  • Use coefficients (numbers before compounds)
  • Never change subscripts
  • Start with metals, then nonmetals, then oxygen/hydrogen

Example: Balance CH₄ + O₂ → CO₂ + H₂O

  1. Balance C: 1 C on left, 1 on right ✓
  2. Balance H: 4 H on left, need 2 H₂O: CH₄ + O₂ → CO₂ + 2H₂O
  3. Balance O: 2 + 4 = 6 on right, need 3 O₂: CH₄ + 2O₂ → CO₂ + 2H₂O

States of Matter in Equations​

Symbols:

  • (s) = Solid
  • (l) = Liquid
  • (g) = Gas
  • (aq) = Aqueous solution

Example: 2H2(g)+O2(g)→2H2O(l)2H_2(g) + O_2(g) → 2H_2O(l)

Types of Reactions​

Combustion​

Combustion - Burning in oxygen

Fuel+O2→Products+heat\text{Fuel} + O_2 → \text{Products} + \text{heat}

Complete combustion:

  • Hydrocarbon + O₂ → CO₂ + H₂O
  • Example: CH₄ + 2O₂ → CO₂ + 2H₂O

Incomplete combustion:

  • Insufficient oxygen
  • Produces CO (carbon monoxide) and soot

Decomposition​

Decomposition - One reactant breaks into two or more products

Compound→Elements/simpler compounds\text{Compound} → \text{Elements/simpler compounds}

Examples:

  • 2H2O2→2H2O+O22H_2O_2 → 2H_2O + O_2 (hydrogen peroxide)
  • 2CuO→2Cu+O22CuO → 2Cu + O_2 (heat)
  • CaCO3→CaO+CO2CaCO_3 → CaO + CO_2 (heat)

Synthesis (Combination)​

Synthesis - Two or more reactants form one product

A+B→AB\text{A} + \text{B} → \text{AB}

Examples:

  • 2Na+Cl2→2NaCl2Na + Cl_2 → 2NaCl
  • C+O2→CO2C + O_2 → CO_2

Displacement (Single Replacement)​

Displacement - More reactive element replaces less reactive

AB+C→AC+B\text{AB} + \text{C} → \text{AC} + \text{B}

Examples:

  • Zn+2HCl→ZnCl2+H2Zn + 2HCl → ZnCl_2 + H_2 (Zn more reactive than H)
  • Cl2+2KBr→2KCl+Br2Cl_2 + 2KBr → 2KCl + Br_2 (Cl more reactive than Br)

Reactivity series: K > Na > Ca > Mg > Zn > Fe > Cu > C > H > O

Neutralization (Acid-Base)​

Acid + Base → Salt + Water

HCl+NaOH→NaCl+H2OHCl + NaOH → NaCl + H_2O

Redox Reactions​

Redox - Reactions involving electron transfer

Oxidation and Reduction​

Oxidation:

  • Loss of electrons
  • Electron removal
  • Increase in oxidation number

Reduction:

  • Gain of electrons
  • Electron addition
  • Decrease in oxidation number

Memory: OILRIG = Oxidation Is Loss, Reduction Is Gain

Oxidation Numbers​

Rules:

  • Element: 0
  • Monatomic ion: Charge of ion
  • Oxygen: Usually -2 (except peroxides -1)
  • Hydrogen: Usually +1 (except hydrides -1)
  • F: Always -1
  • Sum of oxidation numbers = charge

Examples:

  • In H₂O: H = +1, O = -2
  • In ClO⁻: Cl = +1, O = -2
  • In MnO₄⁻: Mn = +7, O = -2

Redox Equations​

Example: Fe + Cu²⁺ → Fe²⁺ + Cu

  • Fe: 0 → +2 (oxidation)
  • Cu: +2 → 0 (reduction)
  • Fe is oxidized (reducing agent)
  • Cu²⁺ is reduced (oxidizing agent)

Precipitation Reactions​

Precipitation - Aqueous ions form insoluble solid

AB(aq)+CD(aq)→AC(s)+BD(aq)AB(aq) + CD(aq) → AC(s) + BD(aq)

Example: AgNO3(aq)+NaCl(aq)→AgCl(s)+NaNO3(aq)AgNO_3(aq) + NaCl(aq) → AgCl(s) + NaNO_3(aq)

Solubility Rules​

Soluble:

  • All alkali salts (Na, K, NH₄)
  • All chlorides (except Ag, Pb)
  • All sulfates (except Ba, Ca, Pb)

Insoluble:

  • Most hydroxides (except alkali metals)
  • Most carbonates (except alkali metals)
  • Most phosphates (except alkali metals)

Ionic Equations​

Ionic equation - Shows only species that change

Full equation: AgNO3(aq)+NaCl(aq)→AgCl(s)+NaNO3(aq)AgNO_3(aq) + NaCl(aq) → AgCl(s) + NaNO_3(aq)

Ionic: Ag+(aq)+Cl−(aq)→AgCl(s)Ag^+(aq) + Cl^-(aq) → AgCl(s)

Spectator ions (Na⁺, NO₃⁻) omitted

Energy in Reactions​

Exothermic - Releases energy (heat)

  • Temperature increases
  • ΔH = negative
  • Combustion, neutralization

Endothermic - Absorbs energy

  • Temperature decreases
  • ΔH = positive
  • Photosynthesis, evaporation

Key Points​

  1. Balance equations using coefficients
  2. Combustion produces CO₂ + H₂O
  3. Four main reaction types
  4. Redox: Electron transfer
  5. Precipitation: Forms insoluble solid
  6. Oxidation numbers track electron movement
  7. Exothermic vs endothermic

Practice Questions​

  1. Balance chemical equations
  2. Classify reaction types
  3. Determine oxidation numbers
  4. Identify oxidation/reduction
  5. Write ionic equations
  6. Predict products from reactants

Revision Tips​

  • Practice balancing equations
  • Know reaction type characteristics
  • Learn oxidation number rules
  • Understand redox clearly
  • Know solubility rules
  • Predict reaction products
  • Practice equation writing