Distinguish between the strength and the concentration of acids and bases, including investigation of the degree of ionisation and the relationship between ionisation, conductivity, and pH

What this dot point is asking

NESA wants you to clearly distinguish two ideas that students often blur: the strength of an acid (its degree of ionisation, an intrinsic property) and the concentration of an acid (how much is dissolved, an extrinsic property). You should be able to compare pH, conductivity and reactivity of strong and weak acids at equal concentration, and explain everything in terms of the position of the ionisation equilibrium. This builds on properties of acids and bases and is the prerequisite for pH calculations and titration curve shapes.

The answer

Strength vs concentration

• Strength describes the degree of ionisation in water. A strong acid ionises essentially 100 percent. A weak acid ionises only a few percent (often less than 1 percent at typical concentrations).
• Concentration describes how much acid is dissolved per litre, regardless of how much has ionised.

These properties are independent: you can have a dilute strong acid or a concentrated weak acid.

Strong acids and bases

Strong acids are essentially fully ionised in water. The common examples for HSC:

• IMATH_5 , $HBr$, $HI$ (hydrohalic acids except $HF$, which is weak).
• IMATH_9 (nitric acid).
• IMATH_10 (sulfuric acid, fully ionised for the first proton, partially for the second).
• IMATH_11 (perchloric acid).

Strong bases are also fully ionised:

• Group 1 hydroxides: $NaOH$, $KOH$.
• Heavier Group 2 hydroxides: $Ca(OH)_2$, $Ba(OH)_2$.

Weak acids and bases

A weak acid only partially ionises, setting up an equilibrium:

The acid dissociation constant ($K_a$) measures the position of this equilibrium:

A small $K_a$ means a weak acid (very little ionised); a large $K_a$ means a stronger acid.

For weak bases the equivalent constant is $K_b$. For ammonia, $K_b = 1.8 \times 10^{-5}$.

Conductivity

Electrical conductivity of a solution depends on the total concentration of ions. At equal acid concentration, a strong acid produces many more ions than a weak acid, so it conducts much better.

This is a definitive HSC experiment: dip a conductivity probe (or a small bulb circuit) into 0.1 mol/L $HCl$ and 0.1 mol/L $CH_3COOH$. The $HCl$ lights the bulb brightly; the ethanoic acid lights it dimly. Both have the same nominal concentration; only the degree of ionisation differs.

pH at equal concentration

For two solutions of equal concentration:

• Strong acid has the lower pH (more ions, lower $[H^+]$ value... wait, higher $[H^+]$, so lower pH).
• Weak acid has a higher pH (less ionised, lower $[H^+]$).

For a 0.10 mol/L solution:

• IMATH_40 : $[H^+] = 0.10$ mol/L, pH = 1.00.
• IMATH_42 ($K_a = 1.8 \times 10^{-5}$): $[H^+] \approx 1.34 \times 10^{-3}$ mol/L, pH = 2.87.

Reactivity at equal concentration

A strong acid reacts faster initially because more $H^+$ is present at any instant. However, a weak acid reacts to the same final extent with a stoichiometric excess of, for example, magnesium, because as $H^+$ is consumed the equilibrium shifts right (Le Chatelier) and more weak acid ionises.

Mark schemes reward students who explicitly separate kinetics (initial rate, set by $[H^+]$) from stoichiometry (total Mg consumed, set by moles of acid).

Dilution effect on degree of ionisation

When a weak acid is diluted, the percent ionisation increases. Mathematically, if $K_a$ is fixed and $c$ decreases, then $x/c = \sqrt{K_a/c}$ grows. In the limit of infinite dilution every weak acid becomes 100 percent ionised. This is a Le Chatelier consequence: dilution favours the side with more particles.

Worked example

A student is given an unknown 0.10 mol/L acid solution. The pH is measured to be 3.40. Is the acid strong or weak? Justify quantitatively.

Step 1: Calculate $[H^+]$.

Step 2: Compare to concentration.

If the acid were strong, $[H^+]$ would equal 0.10 mol/L (pH = 1.00). The measured $[H^+]$ is much smaller, so the acid is weak.

Step 3: Calculate the percent ionisation.

Step 4: Estimate $K_a$.

This is consistent with a weak acid (small $K_a$).

Common traps

Confusing strong with concentrated. "Strong" refers to degree of ionisation; "concentrated" refers to mol/L. Both terms describe the acid, but they describe different things. Concentrated and dilute apply equally to strong and weak.

Stating that a weak acid is "less reactive" without qualification. It is less reactive initially per mole of dissolved acid. Given time and excess reactant, the total chemistry can be the same.

Forgetting that diluting a weak acid increases percent ionisation. Dilution does not turn a weak acid into a strong one, but it does increase the fraction ionised.

Ignoring the second ionisation of $H_2SO_4$. For dilute solutions HSC typically treats both protons as fully ionised. For concentrated solutions the second $K_a$ matters. State your assumption.

Reading too much from low conductivity. Pure water has very low conductivity but is not "weakly acidic" (its $[H^+] = 10^{-7}$ comes from auto-ionisation). Always relate conductivity to total ion concentration, not to acidity alone.

In one sentence

Strength is the degree of ionisation (an intrinsic property: strong acids fully ionise, weak acids only partly ionise, with $K_a$ measuring the extent), while concentration is how much acid is dissolved per litre, so at equal concentration a strong acid has a lower pH, higher conductivity, and faster initial reaction than a weak acid even though both can deliver the same total moles of $H^+$ given enough reaction.