Calculate the pH of a buffer containing 0.150 \, mol L^-1 NH_3 and 0.250 \, mol L^-1 NH_4 Cl. pK_b(NH_3) = 4.75, so pK_a(NH_4^+) = 9.25. [3 marks]

QLDChemistrySyllabus dot point

Topic 1: Chemical equilibrium systems

Describe the composition and action of buffer systems, and explain qualitatively how a buffer resists changes in pH on the addition of small amounts of strong acid or strong base

A focused answer to the QCE Chemistry Unit 3 dot point on buffers. Defines a buffer as a weak acid plus its conjugate base in comparable amounts, walks through how each component consumes added strong acid or base, and applies the reasoning to the carbonic acid/hydrogencarbonate buffer in blood. Includes the buffer-question types that appear in IA1 stimulus.

Generated by Claude Opus 4.810 min answerUpdated 2026-05-22

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What this dot point is asking
The answer
Worked IA1 stimulus
Examples in context
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What this dot point is asking

QCAA wants you to describe what a buffer is (a weak acid plus its conjugate base in comparable amounts, or a weak base plus its conjugate acid), explain qualitatively how it absorbs small additions of strong acid or strong base, and connect the chemistry to a real system such as the bicarbonate buffer in blood. The dot point is qualitative in Unit 3 (the Henderson-Hasselbalch equation is treated explicitly in Unit 4 contexts; here QCAA wants conceptual explanation).

The answer

A buffer is a solution that resists changes in pH when small amounts of strong acid or strong base are added. It contains a weak conjugate acid-base pair in comparable amounts, so the conjugate base absorbs added H+ and the weak acid absorbs added OH-. The buffer fails when either component is exhausted.

Composition

Two equivalent compositions both produce a buffer:

Weak acid plus its conjugate base. For example, ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa).
Weak base plus its conjugate acid. For example, ammonia (NH3) and ammonium chloride (NH4Cl).

Strong acids and strong bases cannot buffer because their conjugates are too weak (or too strong) to consume the opposing perturbation.

The two components must be present in comparable amounts. The pH at which a buffer is most effective is approximately the pKa of the weak acid (or 14 - pKb of the weak base). Beyond about a 10:1 ratio in either direction the buffering capacity is significantly degraded.

How a buffer absorbs added strong acid

Take the ethanoate buffer at equilibrium:

CH_3COOH_{(aq)} \rightleftharpoons CH_3COO^-_{(aq)} + H^+_{(aq)}

Add a small amount of HCl. The added H+ is consumed by the conjugate base:

CH_3COO^-_{(aq)} + H^+_{(aq)} \rightarrow CH_3COOH_{(aq)}

What was added as a strong acid (full H+ release into solution) is now bound in the weakly ionised CH3COOH, which only partly re-ionises. The effective contribution to [H+] in solution is much smaller than the strong acid would have produced in unbuffered water.

The ratio [CH3COO-] / [CH3COOH] shifts slightly toward the acid; pH falls only slightly.

How a buffer absorbs added strong base

Add a small amount of NaOH. The added OH- is consumed by the weak acid:

CH_3COOH_{(aq)} + OH^-_{(aq)} \rightarrow CH_3COO^-_{(aq)} + H_2O_{(l)}

The strong base is converted to water plus more conjugate base. The ratio shifts slightly toward the conjugate base; pH rises only slightly.

Buffer capacity

Buffer capacity is the amount of strong acid or base that can be added before the pH changes significantly. It is determined by:

The absolute amounts of the buffer components. A 1.0 mol/L buffer has more capacity than a 0.1 mol/L buffer at the same ratio.
The ratio of the two components. Most effective when the ratio is close to 1:1. A buffer with very little of one component cannot absorb much of the corresponding perturbation.
The pKa of the weak acid relative to the desired pH. A buffer is most useful within about 1 pH unit of the pKa.

When most of the weak acid has been consumed by additions of base, the buffer cannot absorb further base; subsequent addition causes a sharp pH jump. The same is true when most of the conjugate base has been consumed by additions of acid.

Why salt of weak acid plus its weak acid creates a buffer

A subtle point QCAA examines. Sodium ethanoate provides a high initial concentration of CH3COO- directly. By Le Chatelier, this suppresses the ionisation of the added CH3COOH (common-ion effect): the equilibrium shifts left, leaving CH3COOH largely as undissociated weak acid. So you start with both species at comparable concentrations and both ready to act as absorbers of perturbation.

Adding straight ethanoic acid alone gives almost all CH3COOH and very little CH3COO- (only the small amount produced by self-ionisation), so it is not a buffer.

The bicarbonate buffer in blood

The textbook real-world example. Blood pH must remain at 7.4 plus or minus 0.05 for human survival. The bicarbonate buffer is the primary mechanism.

CO_{2(g)} + H_2O_{(l)} \rightleftharpoons H_2CO_{3(aq)} \rightleftharpoons H^+_{(aq)} + HCO_3^-_{(aq)}

The buffering pair is H2CO3 (weak acid) and HCO3- (conjugate base). Normal arterial blood holds approximately [HCO3-] = 24 mmol/L and [H2CO3] = 1.2 mmol/L.

Response to added acid (e.g. lactic acid during exercise):

HCO_3^-_{(aq)} + H^+_{(aq)} \rightarrow H_2CO_{3(aq)}

Response to added base:

H_2CO_{3(aq)} + OH^-_{(aq)} \rightarrow HCO_3^-_{(aq)} + H_2O_{(l)}

Why the lungs matter. H2CO3 produced from acid neutralisation decomposes to CO2 and water; CO2 is exhaled. Removing CO2 keeps [H2CO3] low so the buffer remains effective. Hyperventilation removes CO2 faster, raising pH; hypoventilation retains CO2, lowering pH.

Why the kidneys matter. Kidneys excrete or retain HCO3- on longer timescales, providing slower-acting compensation when respiratory adjustment alone is insufficient.

This is the QCAA-favoured worked example. Expect at least one IA1 or EA question on the bicarbonate buffer per year.

Other examples worth knowing

Carbonate buffer in seawater. CO2-HCO3- equilibrium buffers ocean pH; ocean acidification reflects strain on this buffer from increased CO2 loading.
Phosphate buffer in cells. H2PO4- / HPO4^2- system buffers intracellular pH around 7.2.
Industrial buffers. Citrate, phosphate and tris (Unit 4 organic context) appear in food chemistry and pharmaceuticals.

Worked IA1 stimulus

Stimulus: a student adds 1.0 mL of 0.10 mol/L HCl to (a) 100 mL of pure water at pH 7.0, (b) 100 mL of a buffer containing 0.10 mol/L CH3COOH and 0.10 mol/L CH3COO- at pH 4.74. Compare the pH change in each case.

Pure water:

Moles H+ added = 1.0 x 10^-4.
[H+] in new mixture = 1.0 x 10^-4 / 0.101 L = 9.9 x 10^-4 mol/L.
pH = 3.00. Change: 4 pH units.

Buffer:

Moles H+ added = 1.0 x 10^-4. Moles CH3COO- = 0.010, moles CH3COOH = 0.010 before mixing.
After: CH3COO- consumed by H+, so CH3COO- = 0.010 - 1.0 x 10^-4 = 0.0099 mol, CH3COOH = 0.0101 mol.
New ratio = 0.0099 / 0.0101 = 0.980. log10(0.980) = -0.009.
pH = pKa + log10(ratio) = 4.74 + (-0.009) = 4.73. Change: 0.01 pH units.

The buffer absorbs the same acid load with a pH change roughly 400 times smaller than pure water. This kind of comparison is exactly what IA1 stimulus rewards.

Common traps

Listing the components incorrectly: A solution of HCl plus NaCl is not a buffer because Cl- is not the conjugate base of a weak acid. The conjugate of a strong acid is too weak to absorb H+.
Forgetting that comparable amounts are needed: A 1:100 ratio is not an effective buffer; the minor component runs out almost immediately.
Claiming the pH does not change at all: It changes slightly. The buffer reduces the change, not eliminates it.
Missing the role of the lung in the bicarbonate buffer: The buffering equilibrium chain extends to CO2 gas exchange; QCAA expects you to bring the lungs (and sometimes the kidneys) into the explanation.
Writing the consumption equations with double arrows: When a strong acid or base reacts with the buffer component, the reaction goes essentially to completion (single arrow). The underlying buffer equilibrium has the double arrow.

Examples in context

Example 1. Blood-pH simulation in Townsville sugar-cane harvesters. Cane harvest workers' blood pH must stay near $7.40$ despite lactic-acid generation during heavy lifting. The carbonic-acid / bicarbonate buffer dominates: $\text{H}_2 \text{CO}_3 \rightleftharpoons \text{H}^+ + \text{HCO}_3^-$ , with $\text{p}K_a = 6.10$ . From Henderson-Hasselbalch, $\text{pH} = \text{p}K_a + \log([\text{HCO}_3^-]/[\text{H}_2 \text{CO}_3])$ . For arterial blood with $[\text{HCO}_3^-] = 24 \, \text{mmol L}^{-1}$ and $[\text{CO}_2 + \text{H}_2 \text{CO}_3] = 1.2 \, \text{mmol L}^{-1}$ , $\text{pH} = 6.10 + \log(20) = 7.40$ . Respiration adjusts $\text{CO}_2$ to fine-tune pH within seconds.

Example 2. Acetate buffer in QCAA IA1 titration scenario. A practical brief asks students to prepare a buffer for cane-mill enzyme assays. Mix $50.0 \, \text{mL}$ of $0.200 \, \text{mol L}^{-1}$ acetic acid ( $\text{p}K_a = 4.76$ ) with $50.0 \, \text{mL}$ of $0.200 \, \text{mol L}^{-1}$ sodium acetate. Resulting pH $= 4.76 + \log(1) = 4.76$ . Adding $1.00 \, \text{mL}$ of $0.100 \, \text{mol L}^{-1}$ HCl ( $0.0001 \, \text{mol}$ H $^+$ ) converts $0.0001 \, \text{mol}$ acetate to acetic acid: new ratio $\approx 0.0099/0.0101 \approx 0.98$ , $\text{pH} = 4.76 + \log(0.98) = 4.75$ . The buffer resists pH change effectively.

Try this

Q1. Define a buffer and explain how it resists pH change on addition of a small amount of strong acid. [3 marks]

Cue. Solution containing weak acid + conjugate base. Added H $^+$ consumed by conjugate base; equilibrium maintains pH.

Q2. Calculate the pH of a buffer containing $0.150 \, \text{mol L}^{-1}$ $\text{NH}_3$ and $0.250 \, \text{mol L}^{-1}$ $\text{NH}_4 \text{Cl}$ . $\text{p}K_b(\text{NH}_3) = 4.75$ , so $\text{p}K_a(\text{NH}_4^+) = 9.25$ . [3 marks]

Cue. $\text{pH} = 9.25 + \log(0.150/0.250) = 9.25 - 0.222 = 9.03$ .

Q3. A buffer is prepared from acetic acid and sodium acetate. (a) Write equations showing buffer action against added strong acid and strong base. (b) Explain why buffer capacity is highest when $[\text{HA}] = [\text{A}^-]$ . (c) Predict effect of dilution by a factor of $10$ . [3+2+2 marks]

Cue. (a) $\text{A}^- + \text{H}^+ \rightarrow \text{HA}$ ; $\text{HA} + \text{OH}^- \rightarrow \text{A}^- + \text{H}_2 \text{O}$ . (b) Equal capacity each direction. (c) pH unchanged (ratio constant), capacity reduced.

Exam-style practice questions

Practice questions written in the style of QCAA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

2023 QCAA-style5 marksA buffer solution is prepared by dissolving 0.10 mol of CH3COOH and 0.10 mol of CH3COONa in 1.0 L of water. (a) Write the equilibrium equation that describes the buffer. (b) Using balanced equations, explain how the buffer responds to (i) the addition of a small amount of HCl, (ii) the addition of a small amount of NaOH. (c) State and explain what happens to the buffer's effectiveness when most of the CH3COOH has been consumed.

Show worked answer →

A 5-mark answer needs the equilibrium, two consumption equations, and the capacity explanation.

(a) Equilibrium.

CH_3COOH_{(aq)} + H_2O_{(l)} \rightleftharpoons CH_3COO^-_{(aq)} + H_3O^+_{(aq)}

(b)(i) Adding HCl (H3O+). The conjugate base CH3COO- consumes the added H3O+:

CH_3COO^-_{(aq)} + H_3O^+_{(aq)} \rightarrow CH_3COOH_{(aq)} + H_2O_{(l)}

The strong acid is replaced by the weak ethanoic acid, so pH falls only slightly.

(b)(ii) Adding NaOH (OH-). The weak acid CH3COOH consumes the added OH-:

CH_3COOH_{(aq)} + OH^-_{(aq)} \rightarrow CH_3COO^-_{(aq)} + H_2O_{(l)}

The strong base is replaced by the weak conjugate base, so pH rises only slightly.

(c) Buffer capacity. When most of the CH3COOH has been consumed by repeated additions of strong base, the buffer can no longer absorb additional OH- without a large pH change. The remaining CH3COO- can still absorb H3O+, but the buffer is asymmetric and the pH will jump sharply on further base addition. The buffer has reached its capacity.

Markers reward correct equilibrium notation, both consumption equations with arrows and species, and explicit reference to capacity tied to running out of one component.

2022 QCAA-style3 marksThe bicarbonate buffer in human blood is described by the equilibrium CO2(g) + H2O(l) <-> H2CO3(aq) <-> H+(aq) + HCO3-(aq). Explain qualitatively how this buffer maintains blood pH near 7.4 when lactic acid is produced during exercise, and identify the role of the lungs in supporting buffer function.

Show worked answer →

A 3-mark answer needs the buffer response, the species shift, and the lung connection.

Buffer response. Lactic acid releases H+ into the blood. The conjugate base HCO3- consumes the added H+:

HCO_3^-_{(aq)} + H^+_{(aq)} \rightarrow H_2CO_{3(aq)}

The strong acid (H+) is converted into the weak acid H2CO3, so blood pH falls only marginally.

Lung support. H2CO3 readily decomposes to CO2 and H2O, and CO2 is exhaled. Removing CO2 pulls the leftward equilibrium back, which keeps H2CO3 concentration low and maintains buffering capacity. Without the lungs continuously removing CO2, accumulating H2CO3 would quickly exhaust the buffer.

Markers reward the consumption equation, the conversion of strong to weak acid, and the explicit lung-CO2 link tied to Le Chatelier shifting the equilibrium chain.