Teacher License – Chemistry Complete Review Guide 2026–2027

4. To balance an equation, make the atoms of each element equal on both sides by changing coefficients only, never the subscripts inside formulas.
5. Main reaction types: combination, decomposition, single displacement, double displacement (including neutralisation and precipitation), and combustion.
6. The limiting reagent is used up first and sets the amount of product; percent yield = (actual yield ÷ theoretical yield) × 100.
7. Boyle's law: P₁V₁ = P₂V₂ at constant T and n. Charles's law: V₁/T₁ = V₂/T₂ at constant P and n. Gay-Lussac's law: P₁/T₁ = P₂/T₂ at constant V and n (T in kelvin).
8. The ideal gas law is PV = nRT, and one mole of an ideal gas occupies 22.4 L at STP (0°C and 1 atm).
9. Dalton's law: total pressure = sum of partial pressures, and partial pressure = mole fraction × total pressure. Graham's law: rate of diffusion is inversely proportional to the square root of molar mass.
10. Molarity \(M = \dfrac{n}{V}\) (V in litres), molality \(m = \dfrac{n_{solute}}{kg_{solvent}}\), and mole fraction of a component = its moles ÷ total moles in the solution.
11. In dilution the moles of solute stay constant: M₁V₁ = M₂V₂. Diluting from 10 M to 0.2 M needs a dilution factor of 50.
12. Solubility: a gas dissolves more as pressure rises (Henry's law) and less as temperature rises; most solids dissolve more when heated; like dissolves like (polar dissolves polar).
13. Colligative properties: ΔT_f = i·K_f·m, ΔT_b = i·K_b·m, and osmotic pressure π = i·M·R·T; for water K_f = 1.86 °C·kg/mol and K_b = 0.512 °C·kg/mol.
14. Arrhenius: an acid gives H⁺ in water and a base gives OH⁻. Brønsted-Lowry: an acid donates a proton and a base accepts one. Lewis: an acid accepts an electron pair and a base donates one.
15. pH = −log[H⁺], pOH = −log[OH⁻], and pH + pOH = 14 at 25°C. A strong acid ionises completely, so 1.0 × 10⁻⁵ M HCl has pH = 5. Strong acids: HCl, HBr, HI, HNO₃, HClO₄ and the first ionisation of H₂SO₄.
16. At the equivalence point \(M_a V_a \times n_b = M_b V_b \times n_a\), where n is the number of protons or hydroxides per formula unit. The pH there is 7 (strong with strong), above 7 (weak acid with strong base), or below 7 (weak base with strong acid).
17. A buffer is a weak acid with its conjugate base (or the reverse) and resists pH change when a small amount of acid or base is added; pH = pKa + log([A⁻]/[HA]).

Common mistakes in this chapter

• ✕ Confusing molarity (mol per litre of solution) with molality (mol per kilogram of solvent). Molality does not change with temperature; molarity does.
• ✕ Forgetting to convert temperature to kelvin (K = °C + 273) or volume to litres when using PV = nRT with R = 0.0821.
• ✕ Calculating the yield from the excess reactant instead of the limiting reagent, or using masses directly without converting to moles.
• ✕ Forgetting the van 't Hoff factor i for electrolytes in colligative problems: a 1 m NaCl solution gives about twice the particles of a 1 m sugar solution, so its freezing-point depression is larger.
• ✕ Forgetting that a diprotic acid such as H₂SO₄ supplies 2 mol of H⁺ per mole of acid in a titration, which gives a concentration that is half or double the correct one.
• ✕ Assuming the equivalence point always means pH = 7. That holds only for a strong acid with a strong base; with a weak acid pH is above 7, and with a weak base it is below 7.
• ✕ Mixing up strong with concentrated: strength means complete ionisation, while concentration means how much solute is present. Ethanoic acid is weak even when concentrated.
• ✕ Assuming a real gas always behaves ideally. Deviation grows at high pressure and low temperature, when molecular volume and intermolecular forces can no longer be ignored.