Two bottles, both labelled 0.1 mol dm–3, both acids. One has a pH of 1 and the other 2.9. Nothing is wrong with the labels — the difference is how many of the molecules actually let go of their protons.
📚 What you need to know
A strong acid dissociates almost completely in water — write it with a single arrow.
A weak acid dissociates only partially — write it with an equilibrium arrow, lying well to the left.
Strong bases: group 1 hydroxides such as NaOH. Weak bases: NH3 and amines.
Strength is not concentration. Strong and weak describe dissociation; dilute and concentrated describe amount.
For a strong monoprotic acid, [H+] = the concentration of the acid, so its pH can be calculated directly.
Distinguish them by pH, electrical conductivity and rate of reaction — at equal concentrations.
Strong means fully split
Put hydrogen chloride in water and essentially every molecule hands its proton to a water molecule. The position of equilibrium lies so far to the right that we do not bother drawing the equilibrium arrow at all.
A strong acid
HCl(aq) → H+(aq) + Cl–(aq)
Ethanoic acid is a different story. At any moment only about one molecule in a hundred has dissociated; the rest are sitting there intact. Here the equilibrium arrow matters, because the reverse reaction genuinely happens.
A weak acid
CH3COOH(aq) ⇌ CH3COO–(aq) + H+(aq)
Count the loose red circles in each beaker. Both hold the same amount of acid, but the left-hand one has released roughly a hundred times more hydrogen ions.
Strength is not concentration
This is the single most confused pair of words in the topic, so it is worth slowing down. Strong and weak tell you what fraction of the acid dissociates — a fixed property of the substance. Concentrated and dilute tell you how much of it is dissolved in each cubic decimetre — entirely up to whoever made the solution.
Compare the two shaded corners. The concentrated weak acid, pH 2.4, is far more acidic than the dilute strong acid at pH 3.0 — so “strong” does not automatically mean “lower pH”.
If you take one sentence from this page, take this one: strength is a property of the substance, concentration is a property of the solution. You can dilute a strong acid until its pH is 6, and it is still a strong acid. Nothing you do with a measuring cylinder changes what fraction of the molecules dissociate.
Bases work the same way
The pattern repeats exactly. A strong base is fully ionised in solution, so the group 1 hydroxides release essentially all of their OH–. A weak base only partly reacts, so ammonia sits at an equilibrium with most of its molecules unreacted.
CH3COOH and other carboxylic acids, H2CO3, HCN, H2S
Bases
NaOH, KOH, Ba(OH)2
NH3, amines such as CH3NH2
Conjugate partner
very weak
moderately strong
Sulfuric acid is diprotic, and only the first proton comes off completely. The second ionisation, HSO4– ⇌ H+ + SO42–, is a weak-acid equilibrium — which is why 1 mol dm–3 H2SO4 does not give 2 mol dm–3 of H+.
Telling them apart in the lab
All three standard tests come back to the same thing: a strong acid puts more H+ into the solution, and everything else follows from that. The critical condition is that you compare solutions of the same concentration.
The same three tests work for bases, with the readings reversed: a strong base gives a higher pH and a higher conductivity than a weak base of the same concentration.
Test
Strong acid
Weak acid
Why
pH
lower, e.g. 1.0
higher, e.g. 2.9
more of the acid has dissociated, so [H+] is higher
Electrical conductivity
high
low
conduction needs mobile ions, and the weak acid has produced far fewer
Rate with Mg or a carbonate
fast, vigorous fizzing
slower fizzing
rate depends on [H+], not on the total acid present
Volume of alkali to neutralise
the same for both
the same for both
neutralisation uses up all the acid, dissociated or not
That last row is the one examiners love. A weak acid reacts more slowly, but it needs exactly the same volume of alkali to neutralise it. As H+ is removed, the equilibrium shifts right and releases more — Le Chatelier again — until every last molecule has given up its proton.
Why some acids are stronger than others
Acid strength comes down to how easily the bond to hydrogen breaks. Go down group 17 and the halogen atom gets larger, so the H–X bond gets longer and weaker. A weaker bond releases its proton more readily.
Acid strength down group 17
HF < HCl < HBr < HI
WORKED EXAMPLE
Calculate the pH of a 0.050 mol dm–3 solution of nitric acid.
Step 1 — is it strong?HNO₃ is a strong monoprotic acid, so it dissociates completely and gives one H⁺ per molecule.[H⁺] = 0.050 mol dm⁻³Step 2 — take the logpH = −log₁₀(0.050)pH = 1.30This shortcut works only for strong acids. For a weak acid you cannot assume [H⁺] equals the concentration — it is far smaller.
WORKED EXAMPLE
Solution X is 1.0 mol dm–3 ethanoic acid, pH 2.4. Solution Y is 1.0 × 10–4 mol dm–3 hydrochloric acid, pH 4.0. Explain how the weaker acid produces the more acidic solution.
What “weaker” tells youEthanoic acid dissociates only slightly, so a small fraction of its molecules release H⁺.But look at how much there isSolution X contains ten thousand times more acid. A small fraction of an enormous amount still beats all of a tiny amount.concentration outweighs strength herepH depends only on [H⁺]. Strength and concentration both feed into it, and either one can dominate.
WORKED EXAMPLE
You are given two unlabelled solutions, both 0.100 mol dm–3: one hydrochloric acid, one ethanoic acid. Describe two experiments that would identify which is which, and state the expected results.
Experiment 1 — measure the pHUse a calibrated pH meter on equal volumes of each.HCl gives about 1.0; CH₃COOH gives about 2.9Experiment 2 — add equal lengths of magnesium ribbonKeep temperature, volume and concentration the same in both.HCl fizzes noticeably fasterA conductivity meter would do just as well. What you must not do is compare different concentrations — the whole test collapses.
💡 Exam tip
Use the right arrow: → for strong, ⇌ for weak. It is often worth a mark by itself.
Say “partially dissociated” or “the equilibrium lies to the left” — not “it does not dissociate”.
For a strong monoprotic acid, [H+] = concentration. Never assume that for a weak one.
When comparing acids, state clearly that the concentrations are equal.
Remember that the volume of alkali needed is the same for a strong and a weak acid of equal concentration.
Explain trends in strength using bond length and bond strength, not just “it is more reactive”.
⚠️ Common mix-up
Using strong and concentrated as synonyms. They describe completely different things.
Saying a weak acid does not dissociate. It does — just not much, and reversibly.
Assuming a strong acid always has the lower pH. Not if it is far more dilute.
Thinking a weak acid needs less alkali to neutralise. It needs exactly the same.
Doubling [H+] for sulfuric acid. Only the first proton comes off completely.
Comparing solutions of different concentrations and drawing a conclusion about strength.
Up next: Neutralisation Reactions — you have met acids and bases separately. Now put them in the same flask and see what is left behind when the protons have finished moving.
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