Acids and Bases

A Brønsted–Lowry acid donates a proton and a base accepts one. Strong means extensive ionization in water, not high concentration. Conjugate acid–base partners differ by one proton.

A Brønsted–Lowry acid donates a proton and a base accepts one. Conjugate partners differ by one proton. A strong acid is extensively ionized in the relevant aqueous conditions, while a weak acid establishes an appreciable equilibrium with its conjugate base. A dilute strong acid and a concentrated weak acid cannot be ranked in pH from the strength labels alone.

Ka quantifies acid dissociation; a larger Ka indicates greater dissociation tendency for the specified equilibrium. For a conjugate acid–base pair in water, KaKb = Kw at the same temperature. Kw changes with temperature, so neutral pH is not universally exactly 7.

For a sufficiently concentrated strong monoprotic acid, the hydronium concentration is approximately the analytical acid concentration. At extreme dilution, water autoionization cannot be neglected. For a weak acid, use an equilibrium calculation and check any approximation.

For HA in water with initial concentration C and small dissociation, Ka ≈ x²C and x ≈ √(KaC). The approximation requires x to be small relative to C. A salt solution may be acidic or basic if an ion reacts with water; a formula containing no written H⁺ or OH^- does not guarantee neutrality.

A buffer contains appreciable amounts of a weak acid and its conjugate base, or the analogous weak-base pair. Added strong base consumes the weak acid; added strong acid consumes the conjugate base. First perform that essentially complete neutralization bookkeeping, then evaluate the remaining equilibrium.

The Henderson–Hasselbalch relationship is useful when both conjugate partners remain in appreciable amounts. Buffer capacity depends on their amounts, while pH depends strongly on their ratio. Diluting both equally can leave the ratio and approximate pH similar while lowering the capacity to absorb additional acid or base.

An acid more readily donates a proton when the resulting conjugate base is relatively stabilized. Electronegativity, bond strength, charge distribution, and resonance can contribute, depending on the comparison. Use a mechanism appropriate to the particular family instead of one universal periodic shortcut.

For related oxyacids, electron-withdrawing atoms and delocalization can stabilize the conjugate base. For binary hydrides down a group, weaker H–element bonds can become especially important. Every claim should identify which structural feature changes and how it affects the conjugate pair.

For a strong monoprotic acid at an ordinary concentration, hydrogen-ion concentration is approximately the acid concentration. For a weak acid, set up its ionization equilibrium. For an acid–base mixture, perform the reaction stoichiometry first to see what remains before choosing an equilibrium expression. Using a weak-acid formula before accounting for added strong base is a common model-selection error.

At 25°C, pH+pOH=14 under the usual aqueous conditions because Kw=1.0×10^-14. At another temperature, Kw and the neutral pH can differ. A pH calculation should identify the species controlling the solution, not merely apply a memorized equation to whichever concentration appears first.

A weak acid and its conjugate base can consume modest additions of strong base or acid, respectively. Track those stoichiometric changes before applying an equilibrium relation. The Henderson–Hasselbalch equation expresses pH in terms of pKa and the conjugate-base-to-acid ratio when the buffer assumptions are appropriate.

Diluting both buffer components by the same factor leaves their ratio approximately unchanged, so pH may change little, but capacity decreases because fewer moles are available per volume. A buffer can be overwhelmed. Near equal acid and conjugate-base amounts, pH is near pKa; this is not the same as claiming every buffer has pH 7.

Before titrant is added, the analyte’s own acid–base equilibrium controls pH. In the buffer region of a weak-acid/strong-base titration, both acid and conjugate base are present. At half-equivalence, their amounts are equal and pH≈pKa. At equivalence, the weak acid has been converted to conjugate base, whose hydrolysis can make the solution basic.

Beyond equivalence, excess strong base generally dominates the pH calculation. The equivalence volume comes from stoichiometry, while the pH at that volume comes from equilibrium. Choose an indicator whose transition range lies in the steep region near the endpoint; a color change far from equivalence produces a systematic error.

An acid–base problem is first a species problem. Identify strong acids and bases that react essentially stoichiometrically under the stated conditions. Count moles, perform that reaction, and inspect what remains. Excess strong acid, excess strong base, a weak acid alone, a conjugate base alone and a buffer require different subsequent models.

Strength describes ionization tendency; concentration describes amount per volume. A concentrated weak acid and a dilute strong acid cannot be ranked by those adjectives alone. For a weak acid, Ka and formal concentration help determine equilibrium hydronium. Water autoionization can become important at very low acid or base concentrations.

Use the temperature-dependent Kw supplied or appropriate to the conditions. The familiar pH+pOH=14 relation assumes the usual 25°C value. Neutrality means equal hydronium and hydroxide concentrations, not an invariant pH of seven at every temperature. Keep logarithms and concentration units consistent.

A buffer contains appreciable amounts of a weak acid and its conjugate base, or the corresponding weak-base pair. Added OH^- consumes HA and forms A^-; added H₃O⁺ consumes A^- and forms HA. This chemical consumption limits the change in free hydronium compared with an unbuffered solution.

After the strong addition reacts, use the remaining conjugate amounts in an appropriate equilibrium relation. Henderson–Hasselbalch summarizes the ratio under its assumptions; it does not replace neutralization stoichiometry. If one partner is exhausted, the original buffer model no longer applies.

Buffers with the same acid/base ratio can have similar pH but different capacity. More total conjugate material can absorb a larger added amount before a major ratio change. Dilution ideally preserves the ratio approximately while reducing capacity per volume; sufficiently extreme dilution can invalidate simplifying assumptions. State which feature is being compared.

Initially, a weak-acid titration contains mostly the acid and its equilibrium products. Before equivalence, partial neutralization creates a conjugate pair. At half equivalence, equal remaining acid and conjugate-base amounts give pH approximately pKa. At equivalence, the conjugate base can hydrolyze water, so pH need not be seven. Beyond equivalence, excess strong base usually controls pH.

Every stage uses the total relevant solution volume when converting remaining moles to concentration. A steep part of a curve explains why a small added volume can create a large pH change and why indicator choice matters. The endpoint is the observed indicator response; equivalence is the stoichiometric condition.

Relative acidity also has a structural basis. Compare how stable the conjugate base is, using the bonding and electron-distribution information supplied. Electronegativity, bond strength, resonance and electron-withdrawing effects can matter in different comparisons. One isolated periodic trend should not override the actual family of compounds being compared.

First identify strong acid/base stoichiometry, weak acid/base equilibrium, buffer behavior, or salt hydrolysis. If strong acid and base are mixed, neutralize them stoichiometrically before applying an equilibrium expression to what remains. A buffer formula applied before neutralization can use the wrong amounts.

For a weak acid alone, Ka relates equilibrium hydrogen ion, conjugate base, and remaining acid. A small-dissociation approximation may simplify the equation, but verify its size. Concentration and acid strength are different: a dilute strong acid and a concentrated weak acid require different descriptions.

At a weak-acid–strong-base equivalence point, the conjugate base can make the solution basic through hydrolysis. At half-equivalence in the appropriate buffer region, equal acid and conjugate-base amounts give pH=pKa. These are distinct points with distinct chemistry.

The acid/base ratio controls pH in the Henderson–Hasselbalch approximation, while total amounts determine how much added strong acid or base can be absorbed before the ratio changes greatly. Two buffers can have the same pH and different capacities.

After adding strong acid, consume conjugate base and form weak acid; after adding strong base, consume weak acid and form conjugate base. Update moles first. If both species occupy the same final volume, their concentration ratio equals their mole ratio, which can simplify the calculation.

An indicator should change color near the steep pH change appropriate to the titration. A chosen endpoint is a practical observation, not identical by definition to stoichiometric equivalence. Explain any systematic offset using the reaction and calculation.