n = C × V
moles = concentration × volume
pH = pKa + log([A⁻]/[HA])
Henderson-Hasselbalch (buffer region)
Acid-base titration is a quantitative analytical technique used to determine the concentration of an unknown acid or base solution by reacting it with a solution of known concentration (the titrant). During the titration, the titrant is gradually added to the analyte until the reaction reaches its equivalence point—the point at which stoichiometrically equivalent amounts of acid and base have reacted.
The pH of the solution changes throughout the titration process, creating a characteristic titration curve. Understanding these pH changes is crucial for selecting appropriate indicators and interpreting results. This calculator helps you determine the pH at any point during a titration based on the volumes and concentrations of the solutions involved.
The pH calculation during titration depends on the type of acid-base reaction and the amount of titrant added. For strong acid-strong base titrations, the pH is determined by the excess H⁺ or OH⁻ ions present after partial neutralization. At the equivalence point of such titrations, the pH is exactly 7.0.
For titrations involving weak acids or bases, the Henderson-Hasselbalch equation becomes essential in the buffer region: pH = pKa + log([A⁻]/[HA]). The equivalence point pH is not 7.0 for weak acid-strong base titrations (pH greater than 7) or strong acid-weak base titrations (pH less than 7) because the conjugate species formed affects the solution pH.
Strong Acid - Strong Base
The most straightforward titration type. The equivalence point pH is exactly 7.0. Examples include HCl with NaOH or H₂SO₄ with KOH. The titration curve shows a sharp vertical rise at the equivalence point.
Weak Acid - Strong Base
Common in laboratory practice. The equivalence point pH is greater than 7 due to the basic conjugate formed. A buffer region exists before the equivalence point. Example: acetic acid with NaOH.
Strong Acid - Weak Base
The equivalence point pH is less than 7 because the conjugate acid of the weak base is formed. Example: HCl with ammonia (NH₃).
Weak Acid - Weak Base
The most complex case with no sharp equivalence point. The pH change is gradual throughout, making endpoint detection difficult. Special indicators or potentiometric methods are often needed.
This calculator assumes ideal solution behavior and complete dissociation of strong acids and bases. In real laboratory conditions, factors such as temperature, ionic strength, and activity coefficients can affect the actual pH values observed. The calculations also assume a 1:1 stoichiometric ratio between the acid and base.
For polyprotic acids (acids that can donate more than one proton), multiple equivalence points exist, and the calculations become more complex. This calculator handles monoprotic acids and bases. For more complex systems, consider the individual dissociation steps and their respective equilibrium constants.
The weak-weak titration calculations provided are simplified approximations. In practice, these titrations are rarely performed due to the poorly defined equivalence points and the difficulty in detecting the endpoint.
Acid-base titrations have numerous practical applications across various fields. In the food industry, titrations determine the acidity of products like vinegar, wine, and fruit juices. Environmental scientists use titrations to measure the alkalinity of water samples and assess acid rain impacts.
In pharmaceutical analysis, titrations verify the purity and concentration of drug formulations. Clinical laboratories use titration techniques for measuring bicarbonate levels in blood samples. Industrial applications include quality control in manufacturing processes for acids, bases, and buffer solutions.
Understanding titration calculations helps students and professionals design experiments, select appropriate indicators, and interpret results accurately. The ability to predict pH changes throughout a titration is fundamental to analytical chemistry education and practice.