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.
Key Formula
Ma × Va = Mb × Vb (for 1:1 stoichiometry)
LaTeX: M_a V_a = M_b V_b
| Symbol | Meaning | Unit |
|---|---|---|
| Ma | Molarity of acid solution | mol/L |
| Va | Volume of acid used | L or mL |
| Mb | Molarity of base solution | mol/L |
| Vb | Volume of base used | L or mL |
Worked Example
Problem
25.0 mL of sodium hydroxide (NaOH) solution of unknown concentration is titrated with 0.100 mol/L hydrochloric acid (HCl). The equivalence point is reached after adding 18.5 mL of HCl. Find the concentration of NaOH.
Solution
Step 1: The neutralisation reaction is: HCl + NaOH → NaCl + H₂O (1:1 molar ratio) Step 2: Use Ma × Va = Mb × Vb Step 3: (0.100 mol/L) × (18.5 mL) = Mb × (25.0 mL) Step 4: Mb = (0.100 × 18.5) / 25.0 = 1.85 / 25.0 = 0.074 mol/L
Answer
[NaOH] = 0.074 mol/L
Types of Acid-Base Titrations and Suitable Indicators
| Titration Type | Example Reaction | pH at Equivalence Point | Suitable Indicator |
|---|---|---|---|
| Strong acid + Strong base | HCl + NaOH | ~7 | Phenolphthalein or methyl orange |
| Weak acid + Strong base | CH₃COOH + NaOH | >7 (~8.7) | Phenolphthalein (pH 8.2–10) |
| Strong acid + Weak base | HCl + NH₃ | <7 (~5.3) | Methyl orange (pH 3.2–4.4) |
| Weak acid + Weak base | CH₃COOH + NH₃ | ~7 (variable) | Litmus or potentiometer |
| Polyprotic acid (1st equiv.) | H₂SO₄ + NaOH (1:1) | Variable | Methyl orange |
| Polyprotic acid (2nd equiv.) | H₂SO₄ + NaOH (1:2) | Variable | Phenolphthalein |
Interactive Tools
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Related Terms
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.
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.
Acid (Chemistry)
An acid is a substance that donates protons (hydrogen ions, H⁺) to another substance or accepts electron pairs, resulting in a sour taste, ability to turn blue litmus red, and a pH below 7 in aqueous solution. Acids play a fundamental role in chemical reactions, biological processes, and industrial applications ranging from digestion to manufacturing. Common examples include hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and acetic acid (CH₃COOH) found in vinegar.
From French "titrer" (to determine the title/standard of), derived from "titre" (title or fineness of gold/silver). In chemistry, it was adopted in the 19th century to describe the process of determining the "strength" or concentration of a solution through standardised reactions.