Acids and Bases Study Guide
Acids and bases show up everywhere — from stomach acid and soap to industrial processes — and the exam version of the topic is mostly about the pH scale and the equilibrium constants behind it. This guide walks the three acid-base models, explains what 'strong' really means versus 'concentrated,' builds pH and pOH fluency, and finishes with Ka/Kb, conjugate pairs, and buffers. The common mistakes section targets the exact errors that cost points: log slips, the strong-vs-concentrated confusion, and sulfuric acid's second proton.
Key formulas
pH = −log[H+], where [H+] is in mol/L
pOH = −log[OH−], where [OH−] is in mol/L
pH + pOH = 14.00 at 25 °C
Kw = [H+][OH−] = 1.0 x 10^−14 at 25 °C
Ka = [H+][A−]/[HA] for the weak acid HA
pKa = −log Ka; pKb = −log Kb
pH = pKa + log([A−]/[HA]) for a buffer (Henderson-Hasselbalch)
Ka x Kb = Kw for a conjugate acid-base pair
Topic checklist
Acid-base models
Arrhenius: acids produce H+ in water, bases produce OH−; limited but historically important.
Brønsted-Lowry: acids are proton (H+) donors, bases are proton acceptors — the working definition for most problems.
Lewis: acids accept electron pairs, bases donate electron pairs — the broadest model.
Strong versus weak
Strong acids (HCl, HBr, HI, HNO3, HClO4, H2SO4 first proton) and strong bases (Group 1 hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2) dissociate completely.
Weak acids and bases reach equilibrium instead; their strength is measured by Ka and Kb.
Strong means fully dissociated; concentrated means lots of it — a dilute strong acid can have a mild pH.
pH calculations
pH = −log[H+]; each pH unit is a tenfold change in [H+].
For strong acids, [H+] ≈ the acid's molarity (watch for diprotic H2SO4's second proton).
For weak acids, use the Ka expression or an ICE table to find [H+].
pH, pOH, and Kw
Kw = [H+][OH−] = 1.0 x 10^−14 at 25 °C links the two scales.
pH + pOH = 14.00 at 25 °C lets you convert between scales.
Neutral at 25 °C means pH 7.00; that number shifts if the temperature changes because Kw changes.
Conjugate pairs and Ka/Kb
Every acid has a conjugate base and vice versa (HA <-> H+ + A−).
Ka x Kb = Kw for a conjugate pair, and pKa + pKb = 14.00 at 25 °C.
Strong acids have very weak conjugate bases; weak acids have stronger conjugate bases.
Buffers
Buffers resist pH change: a weak acid plus its conjugate base in comparable amounts.
pH = pKa + log([base]/[acid]) (Henderson-Hasselbalch) predicts buffer pH directly.
Buffers work best within ±1 pH unit of the pKa.
Common mistakes
Mistake: Assuming pH changes linearly with dilution, e.g. doubling water doubles pH
Fix:
Dilution changes [H+] by the dilution factor, then pH = −log[H+]. Halving the concentration raises pH by only 0.30 (log 2), not by a full unit.
Mistake: Confusing 'strong acid' with 'concentrated acid'
Fix:
'Strong' describes dissociation completeness; 'concentrated' describes amount per liter. 0.001 M HCl is strong but dilute (pH 3); 18 M acetic acid is weak but concentrated.
Mistake: Logarithm errors: dropping the negative sign or inverting 10^−pH
Fix:
Keep track of the negative sign and direction: [H+] = 10^−pH. For pH 3.5, [H+] = 10^−3.5 = 3.2 x 10^−4 M, not 10^3.5.
Mistake: Treating both protons of sulfuric acid as fully dissociated
Fix:
Only the FIRST proton of H2SO4 is strong; the second is weak (Ka2 = 1.2 x 10^−2). For dilute H2SO4, approximate [H+] ≈ 2 x M only when both protons effectively contribute, and check with the Ka2 equilibrium for precision.
Exam tips
Memorize the seven strong acids and the strong bases as a list; recognizing them instantly saves minutes per problem.
Use ph-practice for calculation speed — pH from Ka setups and buffer problems are the most frequent exam items.
For buffer questions, try the Henderson-Hasselbalch form first; fall back to ICE tables only when the ratio is far from 1:1 or concentrations are very low.
Rehearse log math (10^−pH, antilog) on your calculator before the exam so sign errors disappear.
Pair this guide with equilibrium-practice: every weak-acid problem is an equilibrium problem in disguise.