Quantitative analytical technique for determining unknown concentrations through acid-base reactions
Titration is a quantitative analytical technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. The principle behind titration involves the gradual addition of a titrant (the known solution) to a sample until a reaction reaches completion, indicated by a change that can be observed (such as a colour change, pH change, or precipitation).
A laboratory method where a solution of known concentration (titrant) is used to determine the concentration of an unknown solution (analyte). The process involves adding the titrant gradually until the reaction is complete.
Based on stoichiometry – the exact volume of titrant needed to completely react with the analyte is measured. This volume is used to calculate the concentration of the unknown solution using the formula: M₁V₁ = M₂V₂ for 1:1 reactions.
Remember: “Titration finds concentration through controlled reaction.” The key is measuring the exact volume needed for complete reaction.
Formula to remember: M₁V₁ = M₂V₂ (for 1:1 stoichiometry). Always check the mole ratio first!
Definition: The point at which the amount of titrant added is exactly enough to completely neutralize the analyte (the unknown solution).
Definition: The point at which a noticeable change occurs (often indicated by a colour change due to an indicator), which is ideally close to the equivalence point.
Definition: A substance that changes colour at (or near) the equivalence point, used to visually signal the end point of the titration.
Key Difference: Equivalence point is theoretical (moles acid = moles base). End point is experimental (when indicator changes colour).
Choosing an indicator: Select one whose colour change range includes the pH at the equivalence point of your specific titration.
Example: HCl titrated with NaOH
Reaction: HCl + NaOH → NaCl + H₂O
Any indicator with pH range around 4-10 works well due to sharp pH change
Example: CH₃COOH titrated with NaOH
Reaction: CH₃COOH + NaOH → CH₃COONa + H₂O
Must change in basic range (phenolphthalein)
Example: HCl titrated with Na₂CO₃ or NaHCO₃
Reaction: 2HCl + Na₂CO₃ → 2NaCl + CO₂ + H₂O
Must change in acidic range (methyl orange)
Example: CH₃COOH titrated with NH₃
Reaction: CH₃COOH + NH₃ → CH₃COONH₄
Often no suitable visual indicator; use pH meter
Remember:
Determine the exact strength of a prepared HCl solution using titration with 0.1 M NaOH solution.
Titration is based on neutralization: acid + base → salt + water. The point of complete reaction (equivalence point) is detected using an indicator.
Acid-base neutralization: HCl + NaOH → NaCl + H₂O. At equivalence point: moles HCl = moles NaOH. Using M₁V₁ = M₂V₂ (1:1 stoichiometry).
Conc. HCl is approximately 12 M. Prepare 250 cm³ of 0.1 M HCl using dilution formula:
M₁V₁ = M₂V₂ → 12 × V₁ = 0.1 × 250 → V₁ = 2.08 cm³
Transfer 2.08 cm³ conc. HCl to 250 cm³ volumetric flask. Add distilled water to mark and shake well.
Rinse and fill burette with prepared HCl solution. Ensure no air bubbles in tip.
Transfer 10 cm³ of standard 0.1 M NaOH solution to conical flask using pipette.
Add 2-3 drops of phenolphthalein indicator. Solution turns pink (basic).
Slowly add HCl from burette while swirling flask. Continue until pink colour just disappears (colourless or very light pink).
Record initial and final burette readings. Calculate volume of HCl used.
Repeat titration 2-3 times until concordant readings (±0.1 cm³) are obtained.
| Sr. No. | Initial Reading (cm³) | Final Reading (cm³) | Volume of HCl used (cm³) |
|---|---|---|---|
| 1 | 0.0 | 10.0 | 10.0 |
| 2 | 10.0 | 20.0 | 10.0 |
| 3 | 20.0 | 30.0 | 10.0 |
| Mean Volume (V₂): | 10.0 cm³ | ||
Given:
Molarity of NaOH (M₁) = 0.1 M
Volume of NaOH (V₁) = 10 cm³
Volume of HCl (V₂) = 10 cm³ (mean)
Molarity of HCl (M₂) = ?
Reaction: HCl + NaOH → NaCl + H₂O (1:1 mole ratio)
Using formula: M₁V₁/n₁ = M₂V₂/n₂
(0.1 × 10)/1 = (M₂ × 10)/1
1 = M₂ × 10
M₂ = 1/10 = 0.1 M
The exact concentration of HCl solution is 0.1 M.
Answer: The equivalence point is the theoretical point at which the amount of titrant added is exactly stoichiometrically equivalent to the amount of analyte present. It represents complete reaction. The end point is the experimental point where a detectable change (usually colour change of an indicator) occurs. The end point should ideally coincide with the equivalence point, but there may be a slight difference known as indicator error.
Answer: Phenolphthalein changes colour in the pH range of 8.2-10.0 (colourless in acid, pink in base). In weak acid-strong base titrations, the equivalence point occurs at pH > 7 (basic) due to the formation of the conjugate base. Phenolphthalein’s colour change occurs within this basic range. In weak base-strong acid titrations, the equivalence point occurs at pH < 7 (acidic), so phenolphthalein would change colour before the equivalence point, giving inaccurate results. Methyl orange (pH range 3.1-4.4) is more suitable for weak base-strong acid titrations.
Answer: Using the formula M₁V₁ = M₂V₂ (for 1:1 stoichiometry of NaOH and HCl):
M₁V₁ (acid) = M₂V₂ (base)
0.2 M × 25 cm³ = 0.1 M × V₂
5 = 0.1 × V₂
V₂ = 5 / 0.1 = 50 cm³
Therefore, 50 cm³ of 0.1 M NaOH is required to neutralize 25 cm³ of 0.2 M HCl.
Answer: Concordant readings in titration are two or more volume readings that agree within a specified range (usually ±0.1 cm³). They are important because:
Typically, at least two concordant readings are required, and the first rough titration is not included in calculating the mean.
Answer: In weak acid-weak base titrations, both the acid and base are only partially dissociated. At the equivalence point, the salt formed (e.g., ammonium acetate from acetic acid and ammonia) undergoes hydrolysis of both cations and anions, resulting in a solution that is nearly neutral. The hydrolysis reactions buffer the pH change, preventing a sharp rise or fall. Since both conjugate acid and conjugate base are present and can react with added H⁺ or OH⁻ ions, the pH changes gradually throughout the titration. This gradual change makes it difficult to identify the exact equivalence point using visual indicators, which is why potentiometric titration (using a pH meter) is preferred for weak acid-weak base titrations.
Quality control of drugs, determining active ingredient concentrations, testing purity of pharmaceutical products.
Testing water hardness, chlorine content, pH adjustment, monitoring water quality in treatment plants.
Determining acidity of fruit juices, vinegar, wine; testing fat content; quality control in food production.
Quality control of raw materials and final products, concentration determination in chemical processes.
Testing acid rain, monitoring pollution levels, analyzing soil and water samples for contaminants.
Blood tests, cholesterol measurements, glucose monitoring, analysis of body fluids.
Everyday Examples:
Test your understanding of titration concepts with these 20 multiple choice questions. Select your answer and click submit to check your score.