Strength is not concentration
A Brønsted acid donates a proton; a base accepts one. Strong acids dissociate essentially completely in water at ordinary concentrations, while weak acids establish equilibria. Strength describes the tendency to transfer a proton; concentration describes how much substance is present. A dilute strong acid and a concentrated weak acid can have very different pH values.
A proton transfer always has partners. If HA donates H⁺ to water, it forms A⁻ and hydronium; A⁻ is HA's conjugate base. Some substances can act as acid or base depending on the partner. “Strong” in this context refers to how extensively the transfer occurs in water, not to the amount poured, corrosiveness in every setting, or how fast a reaction proceeds.
Water itself participates in acid–base equilibrium. Its ion product links hydronium and hydroxide activities, often approximated by concentrations. At 25 °C, the familiar neutral pH of about 7 follows from equal amounts of the two ions; at another temperature, a neutral solution still has equal hydronium and hydroxide but the numerical pH may change.
A logarithmic scale
pH encodes hydronium activity, often approximated by concentration in introductory problems. One pH unit corresponds to a tenfold change in hydronium concentration in that approximation. Conjugate acid–base pairs differ by one proton; acid and base equilibrium constants are linked through water’s ionization at a given temperature.
Because pH is logarithmic, a change from pH 3 to pH 4 corresponds to about a tenfold decrease in hydronium concentration in dilute idealized water. That is not a tenfold change in the number of total acid molecules unless dissociation and buffering behavior are accounted for. pH can also be measured outside the familiar 0–14 range for sufficiently concentrated solutions; the classroom range is common, not an absolute rule.
Weak acids establish an equilibrium, so dilution can change the fraction ionized even while lowering the overall concentration. A weak acid's pH requires both its amount and its Ka, and possibly other equilibria. If two acids have the same initial concentration, the stronger one generally produces more hydronium; if concentrations differ, strength alone is insufficient to rank pH.
Buffers and titrations
A buffer combines a weak acid with its conjugate base, consuming modest added acid or base. It works best when both components are present in appreciable amounts and has finite capacity. A titration curve traces the changing solution; the pH at equivalence is not always 7, because conjugate species may react with water.
A buffer works because added hydronium can react with the conjugate base and added hydroxide can react with the weak acid. Both components are consumed as the buffer absorbs additions, so its capacity is finite. Near a 1:1 acid-to-base ratio, its pH is near pKa and small additions often produce modest changes. The buffer is not a device that freezes pH at one exact value under any amount of added acid.
During a titration, stoichiometry controls the amount of acid or base left at each stage. Before equivalence, a weak-acid/strong-base mixture can be buffered; at equivalence, the conjugate base can make pH exceed 7 at 25 °C; afterward, excess strong base dominates. The shape of the full curve guides indicator choice and interpretation better than one isolated color.
Comparing weak acids
Ka describes the equilibrium tendency of an acid to donate a proton in water; a larger Ka means a stronger acid in that solvent. Its logarithmic counterpart pKa makes comparisons easier. The Henderson–Hasselbalch relationship estimates buffer pH from pKa and the base-to-acid ratio when its approximations apply. Near equal amounts, pH is close to pKa.
Ka is written for the acid donating a proton to water; Kb describes its conjugate base accepting one. Their product is related to water's ion product at the same temperature. A lower pKa means a stronger acid, because pKa is a negative logarithm of Ka. The relative strengths of a conjugate pair are linked: a very strong acid leaves a very weak conjugate base in water.
The Henderson–Hasselbalch form is useful when a weak acid and conjugate base are both present and their ratio can be estimated reliably. It is less trustworthy when one component is almost depleted, concentrations are extremely low, or activities depart strongly from concentration approximations. Checking the actual species present is more important than recognizing a convenient equation shape.
Follow the whole titration
Before equivalence in a weak-acid/strong-base titration, a buffer region can appear; near equivalence the pH changes sharply. After equivalence, excess strong base usually dominates. The half-equivalence point is useful because the weak acid and conjugate base amounts are equal there. At equivalence, the conjugate base can make the solution basic, so an indicator must suit that particular curve.
At the half-equivalence point of a simple weak-acid/strong-base titration, half the original acid has been converted into its conjugate base. Under suitable assumptions, their amounts are equal and pH is approximately pKa. At equivalence, all original acid has been neutralized in the stoichiometric sense, but the resulting conjugate base can react with water. That is why “equivalence” and “neutral pH” are not synonyms.
An indicator changes color over a range of pH values. Select one whose transition overlaps the steep portion of the particular curve so a small volume error corresponds to the visual endpoint. A pH probe can provide a fuller curve, but calibration and mixing still matter. If the analyte has multiple acidic protons, separate equivalence regions may appear, depending on how distinct its dissociation steps are.