Buffer uses and ocean acidification
Key idea: Dissolved carbon dioxide releases hydrogen ions in seawater. Carbonate buffering limits the pH fall but uses up carbonate ions, so seawater becomes less alkaline and carbonate less available.
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Guided lesson
Guided steps
- Explain why buffers are useful when pH must stay steady.
- Find the acid–base pair and its reactions with added acid or base.
- Predict pH and carbonate changes before following the CO₂ reactions.
- Follow dissolved carbon dioxide through to fewer carbonate ions.
- Explain the reactions linking more carbon dioxide to fewer carbonate ions.
- Find the hydrogen-ion concentration ratio from the two pH readings.
- Explain why pH alone cannot give the carbonate amount consumed.
- Explain why shell formation gets harder while seawater remains alkaline.
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ContinueBlood normally stays near pH 7.4, and seawater is mildly alkaline. Both contain buffers that limit pH changes. Since the Industrial Revolution, however, the oceans have absorbed so much carbon dioxide that their pH has fallen. This lesson looks at some uses of buffers, and at the chemistry of ocean acidification.
Uses of buffers
Buffers are used wherever pH must be kept within a narrow range:
- in the body: the carbonic acid/hydrogencarbonate system, H₂CO₃ / HCO₃⁻, is an important blood buffer. Other buffers, including proteins, also contribute; lungs and kidneys help regulate the system;
- in the laboratory: buffers keep the pH of a reaction mixture steady when the reaction produces or uses small amounts of acid or base;
- for calibrating pH meters: buffer solutions of known pH are used for calibration, because their pH stays steady.
Carbonate and hydrogencarbonate in seawater
Seawater contains dissolved carbon dioxide, carbonic acid, hydrogencarbonate ions and carbonate ions, linked by equilibria:
The HCO₃⁻ / CO₃²⁻ pair acts as a buffer in the same way as any weak acid and its conjugate base.
If acid is added, carbonate ions remove the H⁺:
If base is added, hydrogencarbonate ions remove the OH⁻:
Check your understanding 1
At fixed temperature, extra dissolved CO₂ forms carbonic acid. Predict both changes before reading the causal explanation.
Show answer
Both fall. Carbonic acid supplies H⁺, lowering pH. CO₃²⁻ accepts some of that H⁺ and becomes HCO₃⁻, reducing carbonate concentration.
Ocean acidification
As the amount of carbon dioxide in the atmosphere increases, more dissolves in the oceans. It forms carbonic acid, which ionises to release H⁺ ions, so the pH of seawater falls. The buffer removes much of this H⁺, but by using up carbonate ions:
So the ocean’s pH falls, more slowly than it would without the buffer, and its carbonate ion concentration falls too. Continued carbon dioxide uptake converts more carbonate to hydrogencarbonate, changing the proportions of the buffer components. Buffering limits the change rather than fixing the pH.
Ocean acidification means a fall in the pH of the oceans. It does not mean that seawater has become acidic: global mean surface-ocean pH has fallen from about 8.2 before industrialisation to about 8.1, so it is still alkaline, just less so. These are global averages; individual samples vary. NOAA explains the observed change and carbonate chemistry.
Exam-style question 1
From atmospheric CO₂ to fewer carbonate ions
Problem
Try this before viewing the solution
Show solution step by step
More CO₂ dissolves
1 markMethod
Link the atmosphere to the ocean.Reason
Carbon dioxide dissolves in seawater.Working
Answer
More atmospheric CO₂ means more dissolved CO₂.Carbonic acid
1 markReason
Dissolved carbon dioxide reacts with water.Working
Answer
CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq)More H⁺ ions
1 markReason
Carbonic acid is a weak acid.Working
Answer
H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq), so [H⁺] increases and the pH falls.Fewer carbonate ions
1 markWorking
Answer
The extra H⁺ ions react with carbonate ions: H⁺(aq) + CO₃²⁻(aq) → HCO₃⁻(aq), so [CO₃²⁻] falls.
Self-mark with the mark scheme
Tick each mark point that your answer includes.
Credit the chain from atmospheric CO₂ through carbonic acid and H⁺ to lower pH and fewer carbonate ions.
Reading a pH change
Because [H⁺] = 10^(-pH), the final-to-initial hydrogen-ion ratio is 10^(pHᵢₙᵢₜᵢₐₗ-pH_final). For example, a pH fall of 0.10 gives a factor 10^(0.10) = 1.26, an increase of about 26%. A small difference in pH can represent a substantial concentration change.
Try it yourself 1
Interpret a change in ocean pH
Problem
The pH of a seawater sample falls from 8.20 to 8.05, and its carbonate ion concentration also falls. Calculate the factor by which [H⁺] has increased. Then use the next response to explain whether pH readings determine carbonate consumption and why a working buffer can still change pH.
Try this before viewing the solution
Show the steps Optional working
Hints
Hint 1
acidification
Is pH proportional to hydrogen-ion concentration, or logarithmic?
Hint 2
the buffer
Does a buffer prevent a change in pH, or reduce it?Show solution step by step
Acidification
Method
Calculate the hydrogen-ion concentration ratio.Reason
Each pH is a negative base-10 logarithm of concentration.Working
Answer
10^(-8.05)/10^(-8.20) = 10^(0.15) = 1.41: [H⁺] has increased by about 41%.
The buffer
Method
Identify what the readings leave unknown.Reason
pH measures hydrogen ions remaining in solution, not all the acid absorbed.Working
Answer
The readings do not give the initial carbonate amount, volume or total acid added. They cannot determine the moles of carbonate used up.
The carbonate evidence
Working
Answer
Carbonate removes added hydrogen ions through CO₃²⁻ + H⁺ → HCO₃⁻, changing the buffer ratio. A working buffer limits the pH fall rather than preventing it, so a fall in pH does not by itself show failure.
Quick check
The 0.15 pH fall is a concentration increase of about 41%, rather than 0.15%. The total amount of acid absorbed remains unknown.
Check your understanding 2
Using the sample above, explain whether the pH readings determine moles of carbonate consumed. Explain how a fall in pH and carbonate concentration can accompany a working buffer.
Show answer
pH gives the hydrogen ions remaining, rather than all the acid absorbed. Initial carbonate amount, sample volume and total acid uptake are unknown, so carbonate consumption cannot be calculated. Carbonate removes added H⁺ as HCO₃⁻; this changes the buffer ratio and limits the pH fall rather than preventing it. A buffer can still be working while its pH falls.
Check your understanding 3
Shellfish and corals build their shells and skeletons from calcium carbonate: Ca²⁺(aq) + CO₃²⁻(aq) → CaCO₃(s). Explain why rising atmospheric carbon dioxide makes this harder, even though seawater stays alkaline.
Show answer
More carbon dioxide dissolves and forms carbonic acid, which increases [H⁺]. The H⁺ ions react with carbonate ions to form hydrogencarbonate, so there are fewer carbonate ions available to form calcium carbonate. This does not depend on the seawater becoming acidic: it depends on the carbonate ions being used up.
Common mistakes
- Saying that ocean acidification means the oceans have become acidic. Their pH has fallen but is still above 7.
- Saying that a fall in pH means the ocean buffer has stopped working. A buffer reduces a change; it does not prevent it.
- Leaving out the loss of carbonate ions when explaining the effect on shellfish and corals.
- Treating H₂CO₃ and CO₃²⁻ as a conjugate pair. They differ by two protons.
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Syllabus and review details
- GCE A-Level H1 Chemistry 2027 · 8873-2027
8873 (2027), complete syllabus
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