Acid-base titrations: finding the equivalence point
Learn what an acid-base titration is, how the equivalence point works, what titration curves look like, and how indicators signal the endpoint.
Quick answer
An acid-base titration determines an unknown concentration by slowly adding a standard solution (titrant) from a buret until the equivalence point — the moment moles of acid equal moles of base — signaled by an indicator's color change.
Example: Titrating 25.00 mL of unknown HCl with 0.100 M NaOH takes 25.00 mL to reach equivalence: M_HCl = (0.100 × 25.00)/25.00 = 0.100 M. The pH at equivalence is 7.00 (strong–strong).
What you will learn
- Describe the apparatus and procedure of an acid-base titration.
- Define the equivalence point and distinguish it from the endpoint.
- Sketch titration curves for strong–strong, weak acid–strong base, and strong acid–weak base.
- Select an appropriate indicator for a given titration.
The titration setup
A titration answers 'how much acid (or base) is in this sample?' A standard solution of known concentration — the titrant — is added slowly from a buret into a flask holding the unknown (analyte) plus a few drops of indicator. The titrant is added until the equivalence point: the instant when moles of added base exactly equal moles of acid present (adjusted for stoichiometry).
At that moment, M_acid × V_acid = M_base × V_base (for 1:1 reactions), and the unknown concentration falls out directly. Precision comes from the buret (±0.05 mL) and the sharpness of the endpoint.
Equivalence point versus endpoint
The equivalence point is the theoretical neutralization moment; the endpoint is the observed indicator color change. In a well-designed titration they coincide.
What the pH is at equivalence depends on what remains:
- Strong acid + strong base (HCl + NaOH): only NaCl and water remain → pH 7.00
- Weak acid + strong base (CH₃COOH + NaOH): acetate remains and hydrolyzes → pH > 7
- Strong acid + weak base (HCl + NH₃): NH₄⁺ remains and hydrolyzes → pH < 7
The 'equivalence is pH 7' rule is a special case, not the law.
Reading a titration curve
Plotting pH against titrant volume gives the titration curve, with four recognizable regions:
- Initial: the pH of the analyte alone
- Buffer region: added titrant converts analyte to its conjugate partner — pH changes slowly (Henderson–Hasselbalch territory)
- The steep jump: near equivalence, one drop swings pH by several units — the bigger the jump, the easier the titration
- Past equivalence: pH levels off at the titrant's pH
The half-equivalence point (half the titrant added) is special: [acid] = [base], so pH = pKa — the standard way to measure pKa experimentally.
Indicators and the endpoint
An indicator is a weak acid whose two forms have different colors, changing over a ~2 pH-unit range. The right indicator's range must bracket the steep jump:
- Strong–strong (jump ~pH 4–10): almost any indicator works
- Weak acid–strong base (jump above 7): phenolphthalein (8.2–10.0)
- Strong acid–weak base (jump below 7): methyl orange (3.1–4.4)
Modern labs often replace eyes with pH meters, recording the whole curve and locating equivalence from the steepest slope — but the chemistry is identical.
Common mistakes
Only for strong acid–strong base titrations. Weak acid–strong base gives pH > 7 at equivalence (the conjugate base hydrolyzes); strong acid–weak base gives pH < 7. The pH at equivalence reflects whatever species remain — not a universal 7.
The indicator is a spectator that changes color near the equivalence point. The neutralization happens whether or not indicator is present — the indicator just makes the invisible visible.
Indicators are themselves weak acids/bases and consume tiny amounts of titrant. A few drops suffice; large amounts shift the endpoint and dull the color change.
Real-world applications
Analytical chemistry: determining unknown acid concentrations.
Pharmaceuticals: assaying drug purity.
Food industry: measuring acidity in wine and dairy.
Environmental testing: acid rain and water hardness analysis.
Frequently asked questions
What is the difference between the equivalence point and the endpoint?
The equivalence point is the theoretical moment when moles of titrant exactly neutralize the analyte. The endpoint is the experimental signal — the indicator's color change. A good titration picks an indicator whose endpoint lands on the equivalence point.
How do I find the equivalence point volume?
When moles of acid equal moles of base: for 25.00 mL of 0.100 M HCl titrated with 0.100 M NaOH, that is 25.00 mL of base. In general: M_acid × V_acid = M_base × V_base (for 1:1 stoichiometry).
How do I choose an indicator?
The curve's shape: strong–strong gives a huge vertical jump (any indicator near pH 7 works); weak acid–strong base jumps at pH > 7 (needs phenolphthalein); strong acid–weak base jumps at pH < 7 (needs methyl orange). The indicator's range must bracket the jump.
How do you know the titrant concentration exactly?
Standardization: the titrant's exact concentration is first determined by titrating it against a primary standard (a pure, stable, weighable substance like potassium hydrogen phthalate). Only then is it used on unknowns.