ACIDBASE

Acid-base chemistry concerns the behavior of substances that donate or accept protons or electron pairs during reactions. The three classical frameworks are Arrhenius, Bronsted-Lowry, and Lewis. Arrhenius acids increase H+ (usually described as hydronium, H3O+) in water, while Arrhenius bases increase OH-. Bronsted-Lowry defines acids as proton donors and bases as proton acceptors. Lewis defines acids as electron-pair acceptors and bases as electron-pair donors.

In aqueous solutions, proton transfer defines most acid-base reactions. The pH scale expresses acidity or basicity: pH = -log[H3O+]. The product [H3O+][OH-] equals Kw, about 1×10^-14 at 25°C, making neutral water pH 7. Strong acids and bases dissociate completely, whereas weak acids and bases dissociate only partially, with equilibrium positions described by their dissociation constants.

Acid strength is quantified by Ka for acids and Kb for bases; pKa and pKb are the negative logarithms of these constants. A lower pKa indicates a stronger acid; a higher pKb indicates a stronger base. For conjugate acid-base pairs, the Henderson-Hasselbalch equation gives the pH of buffer solutions: pH = pKa + log([A-]/[HA]). Buffers resist pH changes near their pKa by neutralizing added acid or base.

Applications and phenomena include biological pH regulation, such as the bicarbonate-carbonic acid system maintaining blood pH, environmental impacts of soil and water acidity, and numerous laboratory techniques like acid-base titrations and the use of indicators. Common examples include strong acids (HCl, H2SO4), strong bases (NaOH, KOH), weak acids (acetic acid), and bases (ammonia).