Titration Curves, Indicators & Buffers (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching A-Level Chemistry, and this video is about titration curves, choosing indicators, and buffers.

A titration produces a pH curve, and the flat buffer region on that curve — where a weak acid resists a big pH change as base is added — is a buffer in action, which is why these three topics are covered together.

A titration uses a burette and pipette to find the volume of one solution that exactly reacts with a known volume of another, with an indicator marking the end point. Plotting the pH throughout that titration produces a pH curve — an S-shaped graph whose steep midpoint is the equivalence point, and whose choice of indicator depends on matching the indicator's colour-change range to that steep section. For a weak acid, the curve also has a flat section before the equivalence point, called the buffer region, where the pH barely changes as base is added.

This video covers titrations and pH curves first, then choosing an indicator from the curve, before spending the rest of its time on buffers — what they are, how to calculate their pH, and where they're used.

A titration finds the concentration of an unknown solution by reacting it with a known volume and concentration of another. A known volume of one solution is measured with a volumetric or graduated pipette into a conical flask, the other is added from a burette, and a few drops of indicator mark the end point when it changes colour. For a simple 1 to 1 reaction like hydrochloric acid and sodium hydroxide, the moles of acid equal the moles of base at that point, which is what lets you calculate an unknown concentration.

Plotting pH against the volume of base added during a titration produces a pH curve — always an S-shape, with a steep midpoint called the equivalence point. There are four combinations of strong and weak acids and bases, and each produces a differently shaped curve: a weak acid or weak base gives a shallower, more gradual rise than a strong one, and a weak acid against a strong base shows a flat section — the buffer region — before its equivalence point. That flat section is where this video is heading next.

An indicator is itself a weak acid, HIn, whose undissociated form and conjugate base are different colours, and the colour you see depends on the ratio of the two. Its pKa is the pH at which the colours are balanced, so an indicator is only suitable for a titration if that colour-change range falls entirely within the curve's steep section. Strong acid–strong base titrations can use methyl red or phenolphthalein; weak acid–strong base needs phenolphthalein specifically; and weak acid–weak base titrations have no sharp equivalence point at all, so no indicator works.

A buffer solution resists changes in pH when small amounts of acid or alkali are added. An acidic buffer is made from a weak acid and its conjugate base — ethanoic acid and sodium ethanoate, for example — so there's a reserve supply of both. Add H plus and it reacts with the reserve of ethanoate; add OH minus and it removes H plus, so more ethanoic acid dissociates to replace it. Either way, the ratio of acid to conjugate base barely shifts, so the pH stays almost constant. A basic buffer works the same way with a weak base and its conjugate acid, like ammonia and ammonium chloride.

The pH of a buffer is found from the Ka of its weak acid and the ratio of the acid to its salt: pH equals pKa plus the log of concentration of salt over concentration of acid. This is the Henderson–Hasselbalch equation. Take a buffer of 0.305 moles per decimetre cubed ethanoic acid and 0.520 moles per decimetre cubed sodium ethanoate, with a Ka of 1.74 times 10 to the minus five moles per decimetre cubed. Rearranging the equilibrium expression gives the H plus concentration equal to Ka multiplied by the concentration of acid over salt concentration. This comes out at 1.02 times 10 to the minus five moles per decimetre cubed, and taking the negative log of that gives a pH of 4.99. Diluting a buffer doesn't change its pH at all, because the ratio of salt to acid stays the same.

One of the most important buffers is in your own blood, where hydrogencarbonate ions, HCO3 minus, keep its pH between 7.35 and 7.45. Carbon dioxide from respiration combines with water to form H plus and HCO3 minus in equilibrium: if H plus rises, that equilibrium shifts left to remove it; if Hplus falls, it shifts right to replace it. Either way, the equilibrium keeps the blood's pH constant, which matters, since a drop in blood pH causes acidosis.

Always give a buffer's pH to exactly two decimal places — examiners penalise anything less. And remember that a buffer only copes with small additions of acid or alkali; add too much, and its pH will change significantly.

A titration produces a pH curve, and choosing the right indicator means matching its colour-change range to the curve's steep section. Where that curve has a flat region instead — for a weak acid — that's a buffer in action: a weak acid and its conjugate base holding the pH almost constant, calculated using pH equals pKa plus log of salt concentration over acid concentration. The same buffering principle keeps your blood's pH constant too.

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