An acid–base reaction is a chemical reaction between an acid and a base. In the Brønsted–Lowry framework, it involves transfer of a proton from an acid to a base. In the broader Lewis framework, a base donates an electron pair to an acid, forming a bond. These complementary descriptions encompass aqueous neutralization, proton-transfer equilibria, and reactions that do not involve hydrogen ions. (goldbook.iupac.org)
Definitions and theoretical frameworks
The Arrhenius description identifies acids as substances that increase the concentration of hydrogen ions in water, and bases as substances that increase hydroxide-ion concentration. In aqueous equations, hydrogen ions are commonly represented as hydronium, H₃O⁺, rather than as isolated H⁺. This framework describes many familiar solution reactions but is restricted to aqueous systems. (openstax.org)
Brønsted–Lowry theory defines an acid as a proton donor and a base as a proton acceptor. An acid therefore reacts only when a suitable proton acceptor is present. The general equation is
[ \mathrm{HA+B\rightleftharpoons A^-+BH^+}. ]
HA and A⁻ form one conjugate acid–base pair, while BH⁺ and B form another. Each pair differs by one proton. For example, hydrogen chloride transfers a proton to ammonia, producing ammonium and chloride ions; water need not participate. (openstax.org)
Lewis acid–base theory extends the description to electron-pair acceptance and donation. In
[ \mathrm{BF_3+NH_3\longrightarrow F_3B\leftarrow NH_3}, ]
boron trifluoride accepts the nitrogen lone pair, forming an adduct with a coordinate covalent bond. No proton is transferred. Similar interactions between metal ions and electron-pair-donating ligands underlie much of coordination chemistry. (openstax.org)
Neutralization and reaction equations
Neutralization commonly denotes reaction of an acid with a base to produce water and a salt. Hydrochloric acid and sodium hydroxide provide a standard example:
[ \mathrm{HCl(aq)+NaOH(aq)\longrightarrow NaCl(aq)+H_2O(l)}. ]
Because these reactants are extensively ionized in dilute aqueous solution, the essential change is expressed by the net ionic chemical equation:
[ \mathrm{H_3O^+(aq)+OH^-(aq)\longrightarrow 2H_2O(l)}. ]
Sodium and chloride ions are spectators: they remain in solution without undergoing the proton-transfer step. (openstax.org)
Not every acid–base reaction produces water, and neutralization does not necessarily leave a solution at pH 7. The final acidity depends on the reacting amounts, acid and base strengths, and subsequent equilibria of the products. A weak acid neutralized with a strong base generally leaves a basic solution at the equivalence point because its conjugate base reacts with water. (openstax.org)
Acid strength and equilibrium
Acid strength describes the tendency to donate a proton in a specified solvent; it is not synonymous with concentration. Strong acids react essentially completely with water under ordinary dilute-solution conditions, whereas weak acids establish a measurable chemical equilibrium:
[ \mathrm{HA+H_2O\rightleftharpoons H_3O^++A^-}. ]
The acid dissociation constant, (K_a), quantifies this equilibrium. In the dilute-solution approximation,
[ K_a\approx \frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]c^\circ}, \qquad pK_a=-\log_{10}K_a, ]
where brackets denote equilibrium concentrations and (c^\circ=1\ \mathrm{mol,L^{-1}}). Larger (K_a), or smaller (pK_a), indicates greater acid strength in the same medium. (goldbook.iupac.org)
Proton-transfer equilibria generally favor the weaker acid and weaker base. For a conjugate pair in water, (K_aK_b=K_w): increasing acid strength corresponds to decreasing conjugate-base strength. The solvent also limits observable strength. Acids that react essentially completely with water are “leveled” to hydronium, although their strengths may be distinguishable in another solvent. (openstax.org)
Polyprotic acids donate multiple protons in successive steps, each with its own dissociation constant. The number of transferable protons is not simply the number of hydrogen atoms in a formula: acetic acid contains four hydrogen atoms but behaves as a monoprotic acid in ordinary aqueous acid–base chemistry. (openstax.org)
Water, pH, and amphiprotic behavior
Water can accept or donate a proton. A substance with both capabilities is called amphiprotic. Water’s self-ionization is
[ \mathrm{2H_2O\rightleftharpoons H_3O^++OH^-}. ]
At 25 °C, the dilute-solution concentration product is approximately (10^{-14}) when concentrations are expressed in molar units. A neutral solution has equal hydronium and hydroxide concentrations and a pH near 7; the neutral value changes with temperature. (openstax.org)
pH is formally defined by hydrogen-ion activity:
[ \mathrm{pH}=-\log_{10}a(\mathrm{H^+}). ]
For sufficiently dilute solutions, activity can often be approximated by concentration divided by its standard value. Activity-based treatment becomes important when solution nonideality makes concentration alone an inadequate measure. (goldbook.iupac.org)
Buffers and quantitative analysis
A buffer solution contains appreciable amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. Added acid is consumed by the basic component, while added base is consumed by the acidic component. Buffering capacity is finite and depends on the amounts present. (openstax.org)
The Henderson–Hasselbalch equation expresses the approximate relationship
[ \mathrm{pH}\approx pK_a+ \log_{10}\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}. ]
It is useful when equilibrium concentrations can be adequately represented by the buffer components’ concentrations; extreme dilution or substantial depletion requires fuller equilibrium calculations. (openstax.org)
In analytical chemistry, acid–base titration determines an unknown amount using a reagent of known concentration. The equivalence point is fixed by reaction stoichiometry, whereas the endpoint is an observed signal, such as an indicator color change. Titration curves relate pH to added reagent volume and reveal how acid strength and product equilibria affect the measurement. (openstax.org)