Acids, Bases, pH and Analytical Chemistry
What you must understand
Brønsted–Lowry acids donate protons, bases accept them; conjugate pairs differ by H⁺. Know strong vs weak acids and bases, and amphiprotic species (water, HCO₃⁻). Use pH = −log[H⁺], pOH, and pH + pOH = 14 at 25 °C. Understand Kw, neutralisation, salt hydrolysis (acidic, basic, neutral salts) and buffers (weak acid + conjugate base).
Titration: standard solution, indicator choice, equivalence point, and calculations using c₁V₁ stoichiometry. Know common indicators and their colour changes. Analytical methods: qualitative tests, gravimetric analysis, and instrumental methods (spectrophotometry, chromatography, flame photometry).
Traps
Strong acid does not mean concentrated. Diluting a strong acid by 10 raises pH by 1, not 10. Indicator choice must match the equivalence pH (strong acid–strong base: phenolphthalein or methyl orange; weak acid–strong base: phenolphthalein). Don't forget stoichiometric coefficients in titration calculations.
Rule of thumb
For a strong monoprotic acid, pH = −log c. For a strong base, pOH = −log c, then pH = 14 − pOH.
Worked example
0.01 mol/dm³ HCl: pH = −log(0.01) = 2. 0.01 mol/dm³ NaOH: pOH = 2, pH = 12.