Equivalence Point
The equivalence point in a titration is the stage at which the amount of titrant added is exactly stoichiometrically equivalent to the amount of analyte present, meaning the acid and base have completely reacted with no excess of either. The pH at the equivalence point depends on the nature of the acid and base involved: neutral (pH 7) for strong acid-strong base, above 7 for weak acid-strong base, and below 7 for strong acid-weak base. The equivalence point is distinct from the endpoint, which is the observed colour change of an indicator and may differ slightly due to indicator choice.
Key Formula
n(acid) = n(base) → Ma × Va = Mb × Vb
LaTeX: n_{\text{acid}} = n_{\text{base}} \quad \Rightarrow \quad M_a V_a = M_b V_b
| Symbol | Meaning | Unit |
|---|---|---|
| n_acid | Moles of acid at equivalence | mol |
| n_base | Moles of base at equivalence | mol |
| Ma, Mb | Molar concentrations of acid and base | mol/L |
| Va, Vb | Volumes of acid and base solutions | L |
Worked Example
Problem
In a strong acid-strong base titration, 30.0 mL of 0.150 mol/L NaOH is titrated against HCl. What volume of 0.200 mol/L HCl is needed to reach the equivalence point?
Solution
Step 1: At equivalence point: moles of HCl = moles of NaOH Step 2: n(NaOH) = 0.150 mol/L × 0.0300 L = 0.00450 mol Step 3: n(HCl) = 0.00450 mol Step 4: V(HCl) = n / M = 0.00450 mol / 0.200 mol/L = 0.0225 L = 22.5 mL
Answer
Volume of HCl required = 22.5 mL
Equivalence Point pH for Different Titration Types
| Acid Type | Base Type | Salt Formed | pH at Equivalence Point |
|---|---|---|---|
| Strong (HCl) | Strong (NaOH) | NaCl (neutral) | 7.0 |
| Weak (CH₃COOH) | Strong (NaOH) | CH₃COONa (basic) | ~8.7 |
| Strong (HCl) | Weak (NH₃) | NH₄Cl (acidic) | ~5.3 |
| Weak (HF) | Strong (NaOH) | NaF (basic) | ~8.1 |
| Weak (CH₃COOH) | Weak (NH₃) | CH₃COONH₄ | ~7.0 (variable) |
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Related Terms
Acid-Base Titration
Acid-base titration is a quantitative analytical technique in which a solution of known concentration (titrant) is gradually added to a solution of unknown concentration (analyte) until the reaction reaches its equivalence point, where the moles of acid exactly equal the moles of base. An indicator or pH meter is used to detect the endpoint of the titration, allowing calculation of the unknown concentration. Titration is widely used in medicine, food testing, environmental science, and quality control to determine the concentration of acids or bases in samples.
pH Scale
The pH scale is a logarithmic measure of the hydrogen ion concentration [H⁺] in a solution, ranging from 0 to 14 at 25 °C, where values below 7 indicate acidic conditions, 7 is neutral, and above 7 is basic or alkaline. Introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909, the scale compresses a trillion-fold range of H⁺ concentrations into a convenient 0–14 range. pH measurement is critical in agriculture, biology, medicine, food science, and environmental monitoring.
Buffer Solution
A buffer solution is an aqueous system that resists significant changes in pH when small amounts of acid or base are added, typically composed of a weak acid and its conjugate base (or a weak base and its conjugate acid) in similar concentrations. Buffers operate by consuming added H⁺ or OH⁻ through equilibrium reactions, maintaining the pH within a narrow range. Buffer systems are critical in biological organisms (blood pH 7.35–7.45 maintained by carbonate/bicarbonate), pharmaceuticals, laboratory experiments, and industrial fermentation.
From Latin "aequivalens" (of equal value), composed of "aequi" (equal) and "valere" (to be strong/worth). In chemistry, the term signifies the point where equal chemical "value" (in terms of reactive equivalents) of acid and base have been combined.