Titration Experiments: Procedure and Precautions Guide

Titration is the lab technique for finding the exact concentration of an unknown solution by reacting it with a solution of known concentration. You add the known solution drop by drop from a burette until the reaction just completes, and the volume used tells you the unknown's strength. It is the most precise way school labs measure acids and bases — and the same idea checks medicines, foods, and water.

A good titration needs three things: a standard solution (exactly known concentration), a sharp endpoint, and concordant readings — two results agreeing within 0.1 mL.

Key Concepts

Standard solution: exactly known molarity (mol/L), prepared from an accurately weighed solid.

Equivalence point: the theoretical moment when chemicals are present in exactly the balanced-equation ratio.

Endpoint: the practical moment you see — the indicator's colour change — which should land near the equivalence point.

Indicators: phenolphthalein (colourless to pink, pH 8.2–10; strong acid vs strong base, weak acid vs strong base) and methyl orange (red to yellow, pH 3.1–4.4; strong acid vs weak base).

Concordant readings: titre values agreeing within 0.10 mL, proving repeatability.

Formulas

Titration formula: M₁V₁/n₁ = M₂V₂/n₂, where M is molarity (mol/L), V is volume (mL or L, consistent both sides), and n₁, n₂ are coefficients from the balanced equation. For 1:1 reactions like HCl + NaOH → NaCl + H₂O it simplifies to M₁V₁ = M₂V₂.

Assumptions and limits: the reaction must be fast, complete, with one known stoichiometry; the indicator's range must bracket the equivalence-point pH — phenolphthalein fails for weak-base titrations.

Worked Example

25.0 mL of unknown NaOH is titrated against 0.100 mol/L HCl; the endpoint needs 20.0 mL of acid. Step 1: the 1:1 equation gives M₁V₁ = M₂V₂. Step 2: x × 25.0 = 0.100 × 20.0. Step 3: x = 2.00 ÷ 25.0 = 0.0800 mol/L. Step 4: check via moles — 0.100 × 0.0200 L = 0.00200 mol acid, so 0.00200 mol base in 0.0250 L gives 0.0800 mol/L. Both routes agree: the NaOH is 0.0800 mol/L.

Common Misconceptions

Misconception 1: "Endpoint and equivalence point are the same." Equivalence is the theoretical exact-completion moment; the endpoint is the colour change you see. A good indicator makes them nearly coincide.

Misconception 2: "Any indicator works for any titration." Phenolphthalein misses the true endpoint in a strong-acid-vs-weak-base titration because the equivalence pH is acidic — methyl orange is correct there.

Misconception 3: "One good reading is enough." Repeat until two readings agree within 0.10 mL and average those concordant values.

Real-World Applications

Titrations verify medicine strength in pharmaceutical labs, measure acidity in juices and vinegar, test chlorine and hardness in drinking water, and determine fertilizer nutrient content. Every batch of antacid tablets is verified by titration before sale.

Procedure: rinse the burette with the titrant, remove air bubbles from the tip, pipette the unknown into a conical flask, add 2–3 drops of indicator, add titrant slowly with swirling, go dropwise near the endpoint, record the titre, and repeat for concordance.

Precautions: rinse glassware correctly (burette and pipette with solution, flask with distilled water only), never leave air bubbles in the jet, read the meniscus at eye level, use the indicator sparingly, and control the last 1 mL drop by drop.

Titration finds an unknown concentration from M₁V₁/n₁ = M₂V₂/n₂ using a standard solution, the right indicator for the equivalence pH, and concordant readings — with correct rinsing and no air bubbles.