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8873 / 2027
Theories of Acids and Bases

Lesson 3 of 3

Buffering and oceans

Explain the response and its limits.

A-Level 8873, revised syllabus (2026-2027)

  • I can explain buffers and their uses.Syllabus

    Syllabus 8873, 3(g). Explain buffers and their uses.

Keep an acid and its conjugate base available

A buffer removes small additions of acid or base, rather than preventing all pH change.

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A small amount of sodium hydroxide is added to a buffer of ethanoic acid and sodium ethanoate. What happens to the pH?

  1. It rises sharply, as it would in water
  2. It falls, because the buffer makes more acid
  3. It rises only very slightly
Show the answer

It rises only very slightly

The added OH- reacts with ethanoic acid to form ethanoate and water, so very little free OH- remains.

An acidic buffer contains appreciable amounts of a weak acid and its conjugate base, for example CH3CO2H and dissolved sodium ethanoate. Added H+ is consumed by ethanoate; added OH- is consumed by the weak acid. Because most added strong acid/base is removed, the pH changes much less than it would in water.

Two complementary responses
  1. Add a little acid

    CH3CO2-(aq) + H+(aq) → CH3CO2H(aq). The conjugate base removes the added proton.

  2. Add a little alkali

    CH3CO2H(aq) + OH-(aq) → CH3CO2-(aq) + H2O(l). The weak acid removes hydroxide.

A basic buffer uses a weak base and its conjugate acid, such as NH3/NH4+. NH3 consumes H+; NH4+ consumes OH-. Buffers help keep enzyme conditions suitable and prevent unwanted pH shifts in formulations. Their capacity is finite: a large acid/base addition can use up one component.

Check your understanding

Why is sodium ethanoate solution on its own not an acidic buffer?

  1. It has ethanoate to remove added acid, but no weak acid to remove added alkali.
  2. It is a salt of a strong acid, so its ions do not react with added acid or alkali.
  3. Its solution is neutral, and a buffer must start with a pH below 7 to work.
Show the answer

It has ethanoate to remove added acid, but no weak acid to remove added alkali.

A buffer needs both a weak acid and its conjugate base in useful amounts, so it can resist both acid and alkali.

Exam-style question

Paper 2 Section B style

This question draws on the Mole Concept and Stoichiometry, and Theories of Acids and Bases.

An antacid tablet of mass 1.20 g contains calcium carbonate and an unreactive binder. A student added the tablet to 50.0 cm3 of 1.00 mol dm-3 hydrochloric acid, an excess. When the fizzing stopped, the solution was boiled for a minute, cooled and made up to 250 cm3 with water. A 25.0 cm3 portion needed 34.60 cm3 of 0.100 mol dm-3 sodium hydroxide for neutralisation. Mr(CaCO3) = 100.1; Kw = 1.0 x 10-14 mol2 dm-6.

  1. (a) Define a Bronsted-Lowry acid and a Bronsted-Lowry base. [2]
  2. (b) Write an equation, with state symbols, for the reaction of calcium carbonate with hydrochloric acid. [1]
  3. (c) Calculate the pH of the 1.00 mol dm-3 hydrochloric acid and of the 0.100 mol dm-3 sodium hydroxide. [3]
  4. (d) Calculate the percentage by mass of calcium carbonate in the tablet. [5]
  5. (e) The pH changes sharply from about 3 to about 11 at the end point of this titration. Explain why phenolphthalein (colour change range pH 8.3 to 10.0) is a suitable indicator. [2]
  6. (f) Suggest why the solution was boiled before the titration, and how the percentage in (d) would change if it had not been boiled. [2]
  7. (g) Explain the difference between a strong acid and a concentrated acid. [2]
  8. (h) Blood contains a buffer made of carbonic acid, H2CO3, and hydrogencarbonate ions, HCO3-. Explain, with an equation, how this buffer keeps the pH of blood nearly constant when a small amount of acid enters it. [3]
 [20]

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Apply the same proton bookkeeping to the ocean

Acidification is a fall in pH, even while seawater remains alkaline.

The required ocean pair is CO32-/HCO3-. Carbonate can accept a proton to form hydrogencarbonate; hydrogencarbonate can donate a proton to react with added hydroxide. This is one part of the wider seawater carbonate system.

The ocean pair responds in both directions
AdditionMain buffer reactionRole
AcidCO32-(aq) + H+(aq) → HCO3-(aq)Carbonate removes H+.
AlkaliHCO3-(aq) + OH-(aq) → CO32-(aq) + H2O(l)Hydrogencarbonate removes OH-.

More atmospheric CO2 promotes more dissolved CO2. Its reaction with water can be represented overall as CO2(aq) + H2O(l) ⇌ H+(aq) + HCO3-(aq). The extra H+ is partly taken up by carbonate, increasing hydrogencarbonate and reducing carbonate availability. Buffering limits the change; it cannot cancel a sustained large input, so ocean pH falls.

Worked exampleRead a pH change as a concentration change

A water sample changes from pH 8.2 to 8.0. Has [H+] increased by only 0.2?

Try it first. Then show the worked stepsHide the worked steps
  1. pH is logarithmic, so compare 10-8.0 with 10-8.2.
  2. Their ratio is 100.2 = 1.5849.
  3. Both pH values exceed 7 at 25 degrees C, yet the second sample has more H+.
Answer

[H+] is about 1.58 times as large, a roughly 58% increase. The sample has acidified without becoming acidic in the pH-below-7 sense.

Check your understanding

The ocean is buffered. Why can rising atmospheric CO2 still change its pH?

  1. CO2 does not dissolve in seawater, so it cannot change the pH of the ocean.
  2. A steady CO2 input uses up the buffer's carbonate, so the pH slowly falls.
  3. The buffer fixes the pH exactly, so any change must come from pollution instead.
Show the answer

A steady CO2 input uses up the buffer's carbonate, so the pH slowly falls.

A buffer resists change only within its capacity. Continued CO2 input shifts the carbonate and hydrogencarbonate balance, so the pH drifts down.