A chemical reaction is a process in which one or more substances (reactants) are converted into one or more new substances (products) with different properties. Signs that a chemical reaction has taken place: (1) Change in state — e.g., gas is produced (effervescence), or a precipitate forms. (2) Change in colour — e.g., iron rusting turns reddish-brown. (3) Change in temperature — heat is released (exothermic) or absorbed (endothermic). (4) Change in smell — e.g., food spoiling. Examples: Magnesium burning in air: 2Mg + O₂ → 2MgO (bright white light). Iron and sulphur heating: Fe + S → FeS. Milk souring (lactic acid fermentation). Rusting of iron: 4Fe + 3O₂ + water → 2Fe₂O₃·3H₂O. Law of Conservation of Mass: In a chemical reaction, the total mass of reactants equals the total mass of products — matter is neither created nor destroyed.

A chemical equation is a symbolic representation of a chemical reaction using chemical formulae and symbols. Word equation: Magnesium + Oxygen → Magnesium oxide. Skeletal (unbalanced) equation: Mg + O₂ → MgO. A chemical equation must show: (1) Reactants on the left side, products on the right side, separated by an arrow (→). (2) Physical states: (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous solution. (3) Conditions above/below arrow: heat (△), catalyst, pressure, light. Example with state symbols: 2H₂(g) + O₂(g) → 2H₂O(l). Why equations need to be balanced: To satisfy the Law of Conservation of Mass — atoms of each element must be equal on both sides. An unbalanced equation like H₂ + O₂ → H₂O is incomplete. Always balance by changing coefficients (large numbers before formulae), never by changing subscripts inside formulae.

Balancing a chemical equation means making the number of atoms of each element equal on both sides. Method — Hit and Trial (inspection): Step 1: Write the skeletal equation. Step 2: Count atoms of each element on both sides. Step 3: Add coefficients to balance, starting with the most complex molecule. Step 4: Verify all atoms are balanced. Example: Fe + H₂O → Fe₃O₄ + H₂ (unbalanced) → 3Fe + 4H₂O → Fe₃O₄ + 4H₂ (balanced). Check: Fe: 3=3 ✓, H: 8=8 ✓, O: 4=4 ✓. Example 2: Al + Cl₂ → AlCl₃ → 2Al + 3Cl₂ → 2AlCl₃ (balanced). Important: Coefficients represent the number of moles of each substance. A balanced equation also tells us the molar ratios in which substances react and are produced.

Chemical reactions are classified into 5 main types based on the nature of the change: (1) Combination (Synthesis) Reaction: Two or more substances combine to form a single product. A + B → AB. Examples: C + O₂ → CO₂; CaO + H₂O → Ca(OH)₂ (slaked lime, exothermic); SO₃ + H₂O → H₂SO₄. (2) Decomposition Reaction: A single compound breaks down into two or more simpler substances. AB → A + B. Examples: 2H₂O → 2H₂ + O₂ (electrolysis); 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂ (thermal decomp.); 2AgCl → 2Ag + Cl₂ (photochemical — basis of photography). (3) Displacement (Single) Reaction: A more reactive element displaces a less reactive element from its compound. A + BC → AC + B. Example: Zn + CuSO₄ → ZnSO₄ + Cu (zinc displaces copper). Fe + CuSO₄ → FeSO₄ + Cu. (4) Double Displacement Reaction: Exchange of ions between two compounds; often a precipitate or water is formed. AB + CD → AD + CB. Example: Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl; NaOH + HCl → NaCl + H₂O. (5) Oxidation-Reduction (Redox) Reaction: simultaneous oxidation and reduction. Overlap with other types is common.

Oxidation: A substance is oxidised when it gains oxygen OR loses hydrogen OR loses electrons. Reduction: A substance is reduced when it loses oxygen OR gains hydrogen OR gains electrons. Oxidation and reduction always occur together in a reaction — such reactions are called Redox reactions. Oxidising agent: The substance that causes oxidation (it itself gets reduced). Reducing agent: The substance that causes reduction (it itself gets oxidised). Examples: CuO + H₂ → Cu + H₂O. Here: CuO is reduced (loses O) → it is the oxidising agent. H₂ is oxidised (gains O) → it is the reducing agent. Another example: 2Mg + O₂ → 2MgO. Mg is oxidised (gains O), O₂ is the oxidising agent. In terms of electrons: Na → Na⁺ + e⁻ (oxidation — loss of electron). Cl₂ + 2e⁻ → 2Cl⁻ (reduction — gain of electron). OIL RIG mnemonic: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

Oxidation reactions are important in everyday life. (1) Corrosion: Metals react with oxygen, moisture, and other substances in the environment. Rusting of iron: 4Fe + 3O₂ + xH₂O → 2Fe₂O₃·xH₂O (hydrated iron oxide — brown rust). Corrosion of copper: Cu + CO₂ + H₂O → CuCO₃·Cu(OH)₂ (green patina). Corrosion of silver: 2Ag + H₂S → Ag₂S + H₂ (black tarnish). Prevention of corrosion: painting, oiling/greasing, galvanisation (coating with zinc), electroplating, using stainless steel or alloys. (2) Rancidity: Fats and oils in food are oxidised by oxygen in the air, producing foul smell and taste — this is called rancidity. Rancid food: butter, chips, biscuits going stale. Prevention of rancidity: (a) Store in airtight containers (reduces O₂ contact). (b) Add antioxidants (e.g., BHA, vitamin E). (c) Store in cold (refrigerate). (d) Use nitrogen gas flushing in packaging (chips packet). (e) Store in dark (some oxidation is light-triggered).

Exothermic reactions release heat energy: burning of natural gas (CH₄ + 2O₂ → CO₂ + 2H₂O + heat), respiration (C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy), burning of quicklime with water (CaO + H₂O → Ca(OH)₂ + heat), decomposition of vegetable matter into compost. Endothermic reactions absorb heat: photosynthesis, decomposition of calcium carbonate (CaCO₃ → CaO + CO₂), electrolysis of water. Heat released or absorbed is shown on the products or reactants side of the equation.

In a combination reaction two or more substances combine to form a single product. Examples: (1) CaO + H₂O → Ca(OH)₂ + heat (slaked lime; used for whitewashing; the surface forms CaCO₃ giving a shiny finish after 2–3 days). (2) 2Mg + O₂ → 2MgO (magnesium burns with a dazzling white flame). (3) C + O₂ → CO₂. (4) 2H₂ + O₂ → 2H₂O. Most combination reactions are exothermic.

A single reactant breaks down into two or more simpler products. Thermal decomposition (heat): CaCO₃ → CaO + CO₂; 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃; 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂ (brown fumes). Electrolytic decomposition: 2H₂O → 2H₂ + O₂. Photolytic decomposition (light): 2AgCl → 2Ag + Cl₂ (white silver chloride turns grey in sunlight; used in black and white photography). Decomposition reactions are the opposite of combination reactions and are mostly endothermic.

A more reactive element displaces a less reactive element from its compound. Examples: Fe + CuSO₄ → FeSO₄ + Cu (blue solution turns pale green; reddish-brown copper deposits on iron); Zn + CuSO₄ → ZnSO₄ + Cu; Zn + 2HCl → ZnCl₂ + H₂. Reactivity order (high to low): K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au. A metal placed above another in the series displaces it from its salt solution; copper cannot displace zinc.

In a double displacement reaction two compounds exchange ions to form two new compounds. Examples: Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl (white precipitate of barium sulphate); AgNO₃ + NaCl → AgCl↓ + NaNO₃ (white precipitate); Pb(NO₃)₂ + 2KI → PbI₂↓ + 2KNO₃ (yellow precipitate). A reaction that produces an insoluble solid (precipitate) is a precipitation reaction. Neutralisation (acid + base → salt + water) is also a double displacement reaction.

According to the Law of Conservation of Mass, the number of atoms of each element must be equal on both sides. Steps: (1) Write the skeletal equation with correct formulae. (2) List the atoms on each side. (3) Balance the compound with the most atoms first by changing coefficients (never change subscripts). (4) Balance hydrogen and oxygen last. (5) Check all atoms. (6) Add state symbols (s), (l), (g), (aq). Example: Fe + H₂O → Fe₃O₄ + H₂ balances to 3Fe + 4H₂O → Fe₃O₄ + 4H₂.

A chemical reaction may be recognised by: change in state, change in colour, evolution of a gas, formation of a precipitate, or change in temperature. Examples: lime water turns milky when CO₂ passes through it (Ca(OH)₂ + CO₂ → CaCO₃↓ + H₂O); magnesium burns with a bright white flame to give white MgO powder; the blue colour of CuSO₄ fades when an iron nail is dipped in it. In an equation, conditions such as heat, light or a catalyst are written above or below the arrow.

(1) Burning magnesium ribbon (cleaned with sandpaper to remove oxide layer): brilliant white flame, white powder MgO. (2) Heating lead nitrate in a boiling tube: brown fumes of nitrogen dioxide, yellow residue of lead oxide. (3) Heating ferrous sulphate crystals: green colour changes to brown, smell of burning sulphur from SO₂ and SO₃. (4) Zinc granules with dilute HCl or H₂SO₄: bubbles of hydrogen gas, test tube becomes warm. (5) Heating copper powder in air: brown copper turns black (CuO).

Types: combination, decomposition, displacement, double displacement, and oxidation–reduction (redox). Oxidation is gain of oxygen or loss of hydrogen (loss of electrons); reduction is loss of oxygen or gain of hydrogen. Key equations: 2Mg + O₂ → 2MgO; CaO + H₂O → Ca(OH)₂; CaCO₃ → CaO + CO₂; 2H₂O → 2H₂ + O₂ (ratio of H₂ to O₂ is 2:1); Zn + CuSO₄ → ZnSO₄ + Cu; CuO + H₂ → Cu + H₂O. Corrosion is the slow oxidation of metals (rusting of iron); rancidity is the oxidation of fats and oils; prevented by painting, galvanisation, antioxidants and nitrogen packing.