What Is Titration?
Titration is a fundamental volumetric technique in which a solution of known concentration (the titrant) is carefully added to a solution of unknown concentration (the analyte) until the reaction between them is complete. The point at which the reaction is just complete is called the equivalence point. By measuring the volume of titrant used, the concentration of the analyte can be calculated precisely.
Titrations are among the most accurate wet-chemical methods available, capable of precisions better than 0.1% when performed carefully with a calibrated burette.
Equipment and Setup
A standard titration uses: - A burette (typically 50 mL, graduated to 0.05 mL) to deliver titrant with precision - An Erlenmeyer flask or beaker to hold the analyte solution - A pipette to transfer a precise volume of analyte - An indicator or pH meter to detect the endpoint - A white tile behind the flask to observe color changes clearly
The endpoint is the experimentally observed point (often a color change) used as a practical approximation of the true equivalence point. The small difference between endpoint and equivalence point is the titration error.
Acid-Base Titrations
Acid-base titrations are the most common type. A strong acid is titrated against a strong base (or vice versa), with the neutralization reaction:
H⁺ + OH⁻ → H₂O
Indicators
Indicators are weak acids or bases that change color at different pH values. Common choices: - Phenolphthalein: colorless below pH 8.2, pink above pH 10. Ideal for strong acid / strong base titrations (equivalence point near pH 7) and weak acid / strong base (equivalence point slightly above 7). - Methyl orange: red below pH 3.1, orange/yellow above pH 4.4. Better for strong acid / weak base systems. - Universal indicator: a mixture that changes through a spectrum of colors across the full pH range.
Calculations
For a monoprotic acid-base system at the equivalence point:
moles of acid = moles of base
n_acid = C_acid × V_acid
n_base = C_base × V_base
Therefore: C_acid × V_acid = C_base × V_base
Example: 25.00 mL of HCl is titrated with 0.1000 mol/L NaOH, requiring 22.45 mL to reach the endpoint.
C_HCl = (0.1000 mol/L × 22.45 mL) / 25.00 mL = 0.08980 mol/L
For polyprotic acids (e.g., H₂SO₄, H₃PO₄), stoichiometric ratios must be applied. H₂SO₄ provides two H⁺ ions per molecule, so 1 mol H₂SO₄ neutralizes 2 mol NaOH.
Back Titration
When an analyte does not react directly or quickly with the titrant, a back titration is used. A known excess of reagent is added to the sample, the reaction is allowed to proceed, and the remaining excess is titrated. The amount of analyte is calculated by difference.
Redox Titrations
Redox titrations exploit electron transfer reactions. One reactant is oxidized while the other is reduced. Common systems include:
- Permanganate titrations: KMnO₄ (deep purple) is reduced to Mn²⁺ (nearly colorless) in acidic solution. MnO₄⁻ acts as its own indicator — the endpoint is signaled by the first permanent faint pink color.
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
- Iodometric titrations: Iodine (I₂) released in a reaction is titrated with sodium thiosulfate (Na₂S₂O₃) using starch as the indicator. The blue-black starch-iodine complex disappears sharply at the endpoint.
I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻
- Dichromate titrations: K₂Cr₂O₇ titrates iron(II) in steel analysis. Diphenylamine sulfonate indicator provides a sharp color change.
Complexometric Titrations
Complexometric titrations use a complexing agent to form a stable complex with metal ions. EDTA (ethylenediaminetetraacetic acid, H₄Y) is the dominant titrant because it forms 1:1 complexes with nearly all metal ions regardless of charge:
M^n+ + Y⁴⁻ → MY^(n-4)
Water Hardness Determination
The classic application is measuring water hardness — the total concentration of Ca²⁺ and Mg²⁺. The sample is buffered to pH 10 with ammonia/ammonium chloride buffer. Eriochrome Black T (EBT) indicator turns wine-red when complexed with metal ions and blue when free. As EDTA is added, it steals metal ions from the indicator complex, and the solution turns blue at the endpoint.
Total hardness (as mg/L CaCO₃) = (V_EDTA × C_EDTA × M_CaCO₃ × 1000) / V_sample
Titration Curves
A titration curve plots pH (for acid-base) or electrode potential (for redox) against volume of titrant added. Key features: - A gradual initial change as titrant is added - A steep inflection point at the equivalence point — this is where pH or potential changes most rapidly - A gradual approach to the titrant's own pH or potential after the equivalence point
The steeper and more pronounced the inflection, the sharper the endpoint detection. Strong acid / strong base systems produce the sharpest curves; weak acid / weak base systems can be too gradual for accurate titration.
Standardization of Solutions
Titrant solutions must be standardized — their exact concentration verified against a primary standard. Primary standards are pure, stable, non-hygroscopic compounds of known molecular weight. Examples: - Potassium hydrogen phthalate (KHP): used to standardize NaOH solutions - Sodium carbonate (Na₂CO₃): used to standardize HCl solutions - Potassium dichromate (K₂Cr₂O₇): used in redox standardization
Applications
Titrations are workhorses of industrial and environmental analysis: - Pharmaceutical manufacturing: Assaying active ingredient content in drug formulations - Food industry: Acidity titration in orange juice, wine, and vinegar; iodometric determination of vitamin C - Water treatment: Total hardness, alkalinity, and dissolved oxygen (Winkler method) measurements - Steel industry: Iron content in ore by permanganate or dichromate titration - Environmental labs: Chloride in wastewater by argentometric (AgNO₃) titration
Mastering titration — from precise burette technique to choosing the right indicator and performing stoichiometric calculations — builds the quantitative reasoning skills central to all of analytical chemistry.