What is the conjugate acid-base pair concept?

Every Bronsted-Lowry acid (proton donor) produces a conjugate base by losing its proton. Every base (proton acceptor) produces a conjugate acid. In $HCl + H_2O \rightarrow H_3O^+ + Cl^-$ the pairs are $HCl / Cl^-$ and $H_2O / H_3O^+$. Strong acids have negligibly weak conjugate bases.

How does the auto-ionisation of water set pH and pOH?

Pure water self-ionises so that $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ at 25 degrees C. Taking $-\log_{10}$ of both sides gives $pH + pOH = 14$. Neutral water has $[H^+] = [OH^-] = 10^{-7}$ so pH equals 7 only at 25 degrees C; at higher temperatures Kw is larger and neutral pH falls below 7.

How do I calculate the pH of a weak acid?

Use an ICE table with $K_a = [H^+][A^-] / [HA]$. For initial concentration $c$ much greater than $K_a$, the approximation $[H^+] = \sqrt{K_a c}$ holds. Always check that x is less than 5 percent of c; otherwise solve the quadratic.

How are buffers designed?

Choose a weak acid whose $pK_a$ is within plus or minus 1 of the target pH, then use the Henderson-Hasselbalch equation $pH = pK_a + \log_{10}([A^-]/[HA])$ to set the ratio of conjugate base to acid. Buffer capacity is maximised when the ratio is 1, i.e. $pH = pK_a$.

How do I choose a titration indicator?

Pick an indicator whose colour-change pH range straddles the equivalence pH. Strong-strong: bromothymol blue (pH 6 to 8). Weak acid plus strong base: phenolphthalein (pH 8 to 10). Strong acid plus weak base: methyl orange (pH 3 to 5). The wrong indicator changes colour at the wrong moles and gives a systematic error.

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NSWChemistry

HSC Chemistry Module 6 Acid/Base Reactions: deep-dive 2026 guide

Deep-dive on HSC Chemistry Module 6 Acid/Base Reactions. Bronsted-Lowry theory, pH and pKa, weak-acid ICE calculations, buffer design, titration curves, and indicator selection at HSC depth.

Generated by Claude OpusReviewed by Better Tuition Academy9 min readNESA-CHEM-MOD-6Updated 2026-05-19

Why Module 6 is the calculation backbone of HSC Chemistry

Module 6 sits at the centre of the HSC syllabus because every other module (organic, electrochemistry, analytical) eventually returns to acid-base equilibria for buffers, pH-dependent reactions, or titration analysis. NESA examiners reward students who can move fluently between conceptual definitions and numerical calculations.

Bronsted-Lowry theory and conjugate pairs

A Bronsted-Lowry acid donates a proton; a base accepts one. The pair before and after proton transfer is the conjugate pair. Amphiprotic species (water, bicarbonate, dihydrogen phosphate) act as either.

Recognising the conjugate pair is the first step in any acid-base question. Examiners often disguise this by giving an organic acid or weak base structure; identify the labile proton.

The pH scale, Kw and temperature

At 25 degrees C $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ . Taking logs gives $pH + pOH = 14$ .

Two HSC traps:

Kw grows with temperature because auto-ionisation is endothermic, so neutral pH falls below 7 above 25 degrees C even though the solution is still neutral.
Strong-acid pH calculations need a sanity check: 0.10 M HCl has pH 1, but 0.10 M H2SO4 (a diprotic strong acid) has pH close to 0.7 because both protons dissociate.

Strong versus weak: Ka and dissociation

Strong acids and bases fully dissociate; weak ones partially dissociate, governed by the dissociation constant $K_a$ or $K_b$ . The relationship $K_a K_b = K_w$ links the strength of an acid and its conjugate base.

For a 0.10 M weak acid HA with $K_a = 1.8 \times 10^{-5}$ :

K_a = \frac{x^2}{0.10 - x} \approx \frac{x^2}{0.10}

x = [H^+] = \sqrt{1.8 \times 10^{-6}} = 1.34 \times 10^{-3} \text{ M}, \quad pH = 2.87

Check the approximation: x is 1.3 percent of 0.10, so the approximation is valid.

Buffers and the Henderson-Hasselbalch equation

A buffer is a solution of a weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change when small amounts of acid or base are added.

pH = pK_a + \log_{10}\left(\frac{[A^-]}{[HA]}\right)

Practical design: to make a buffer at pH 5.0, choose acetic acid ( $pK_a = 4.74$ ) and adjust the ratio. The buffer is effective within plus or minus 1 pH unit of $pK_a$ .

Worked example: blood is buffered by the carbonic acid / bicarbonate system. $pK_a$ of $H_2CO_3$ is about 6.4, and blood pH is 7.4, so $[HCO_3^-] / [H_2CO_3]$ is approximately $10^{1.0}$ , that is 10 to 1. Respiratory and renal compensation adjust both species.

Titrations and titration curves

A titration delivers titrant of known concentration to analyte until the equivalence point, when moles of acid equal moles of base.

Four shapes to recognise:

Strong acid plus strong base: equivalence at pH 7, steep transition.
Weak acid plus strong base: equivalence above pH 7 (the conjugate base is basic). At half-equivalence, $pH = pK_a$ .
Weak base plus strong acid: equivalence below pH 7.
Polyprotic acid: multiple equivalence points (sulfuric, phosphoric, carbonic).

Half-equivalence in case 2 is gold for examiners: you can read $pK_a$ directly off the graph.

Worked example: NaOH into acetic acid

25.00 mL of 0.100 M acetic acid is titrated with 0.100 M NaOH. Calculate pH at four points.

Initial: weak acid only. $pH = 2.87$ (above).
After 12.50 mL NaOH (half-equivalence): equal moles of HA and $A^-$ , $pH = pK_a = 4.74$ .
At equivalence (25.00 mL): all converted to $A^-$ in 50.00 mL total. $[A^-] = 0.0500$ M. $K_b = K_w / K_a = 5.6 \times 10^{-10}$ . $[OH^-] = \sqrt{K_b c} = 5.3 \times 10^{-6}$ . $pOH = 5.28$ , $pH = 8.72$ .
After 35.00 mL NaOH (excess base): excess $OH^-$ in 60.00 mL. $[OH^-] = (0.100 \times 0.010) / 0.0600 = 0.0167$ M, $pH = 12.22$ .

Indicator selection

Indicators are themselves weak acids: $HIn \rightleftharpoons H^+ + In^-$ . Colour change occurs over $pK_{in} \pm 1$ .

Match indicator pH range to the equivalence pH:

Bromothymol blue (6.0 to 7.6): strong-strong.
Phenolphthalein (8.2 to 10.0): weak acid plus strong base.
Methyl orange (3.1 to 4.4): weak base plus strong acid.

Using phenolphthalein for a HCl-NaOH titration would still work but adds a small systematic error; using methyl orange for acetic acid-NaOH gives a large error because methyl orange changes well before equivalence.

Common HSC examiner trap list

Forgetting Kw is temperature-dependent.
Reporting pH to more sig fig than the input concentrations support (the digits after the decimal in pH are the significant ones).
Treating a weak acid as if 100 percent dissociated.
Picking the wrong indicator for the conjugate-pair behaviour at equivalence.
Confusing equivalence (stoichiometric) with end-point (indicator colour change).

In one sentence

Module 6 rewards systematic acid-base reasoning: identify the conjugate pair, classify strong or weak, set up an ICE table or Henderson-Hasselbalch line as appropriate, choose the indicator that matches the equivalence pH, and report answers with correct sig fig.