Introduction
Acid–base chemistry is one of the most frequently applied frameworks in all of chemistry, from biological buffer systems to industrial processes, and builds directly on the equilibrium concepts introduced in Thermodynamics and Kinetics.
1. Arrhenius, Brønsted–Lowry, and Lewis concepts
Three progressively broader definitions of acids and bases are used depending on context. The Arrhenius definition (an acid produces H⁺ in water, a base produces OH⁻) works only for aqueous solutions. The Brønsted–Lowry definition (an acid is a proton donor, a base is a proton acceptor) is broader, covering non-aqueous and gas-phase reactions. The Lewis definition (an acid is an electron-pair acceptor, a base is an electron-pair donor) is the broadest, encompassing reactions with no protons involved at all — such as the formation of coordinate bonds in complex ions (Chapter 8, Chapter 10).
2. Conjugate acid–base pairs
In a Brønsted–Lowry reaction, an acid becomes its conjugate base after donating a proton, and a base becomes its conjugate acid after accepting one. A strong acid has a very weak conjugate base (and vice versa) — this inverse relationship is essential for predicting the position of acid–base equilibria and for choosing appropriate buffer components.
3. Strength, dissociation, and equilibrium
Strong acids/bases dissociate essentially completely in water; weak acids/bases establish an equilibrium between dissociated and undissociated forms, quantified by the acid dissociation constant, Ka (or base constant, Kb). A larger Ka means a stronger (more dissociated) weak acid. Ka and Kb for a conjugate pair are related through the water autoionization constant, Kw = Ka × Kb.
4. pH and related logarithmic measures
pH = −log[H⁺] compresses the enormous range of hydrogen ion concentrations found in real solutions into a manageable scale; each unit change represents a tenfold change in [H⁺]. Related measures include pOH (= −log[OH⁻], with pH + pOH = 14 at 25°C) and pKa (= −log Ka, useful because pH = pKa when a weak acid is exactly half-neutralised — the basis of the Henderson–Hasselbalch equation below).
5. Buffers, hydrolysis, and titration
A buffer resists pH change upon addition of small amounts of acid or base, typically made from a weak acid and its conjugate base (or weak base and its conjugate acid) in comparable amounts; its pH is given by the Henderson–Hasselbalch equation, pH = pKa + log([conjugate base]/[weak acid]). Hydrolysis is the reaction of an ion with water that shifts pH away from neutral — for example, the conjugate base of a weak acid reacts with water to produce a mildly basic solution, which is why salts of weak acids are not pH-neutral. A titration curve (pH plotted against volume of titrant added) shows a gradual pH change during buffering regions and a sharp jump at the equivalence point, where moles of acid and base added are stoichiometrically equal — the shape and position of this jump differs characteristically for strong–strong, strong–weak, and weak–weak acid–base titrations.
Key Takeaways
- Three acid–base definitions (Arrhenius, Brønsted–Lowry, Lewis) apply in progressively broader contexts, with Lewis being general enough to describe coordinate bond formation.
- Strong acids/bases dissociate completely; weak ones establish an equilibrium described by Ka/Kb, with a strong-acid/weak-conjugate-base inverse relationship.
- pH, pOH, and pKa compress concentration ranges logarithmically, with pH = pKa marking the half-neutralisation point of a weak acid.
- Buffers (via the Henderson–Hasselbalch equation) resist pH change; titration curves reveal the equivalence point and distinguish strong versus weak acid–base systems by their shape.