Lesson 6 · 20 min

What makes an acid strong

Why is HI stronger than HF, HClO₄ stronger than HOCl, and an Fe³⁺ solution acidic? Bond energy, electronegativity, and charge delocalisation turn dozens of pKa values into a few rules.

In this lesson

Explain the strength of binary acids with H–A bond energy and the electronegativity of A
Explain oxoacid strength with terminal oxygens and charge delocalisation
Explain the inductive effect and the acidity of hydrated metal ions
Understand water’s levelling effect and meet Lewis acids

Hydrofluoric acid is “weak” yet etches glass; perchloric acid is among the strongest common acids; and a bottle of iron(III) chloride solution can sit near pH 2. Acid strength is not arbitrary — it is written in molecular structure. This lesson threads dozens of $\mathrm{p}K_\mathrm{a}$ values onto a few rules.

Predict first

Among $\ce{HOCl}$ , $\ce{HClO2}$ , $\ce{HClO3}$ , and $\ce{HClO4}$ , which is the strongest acid?

Pick an answer and lock it in before you look. A wrong guess teaches more than no guess.

\ce{HOCl}

\ce{HClO2}

\ce{HClO3}

\ce{HClO4}

Binary acids: bond energy and electronegativity

For acids of the form H–A two factors compete:

Across a period, A becomes more electronegative and the H–A bond more polar, so the proton leaves more easily: $\ce{CH4} \ll \ce{NH3} \ll \ce{H2O} < \ce{HF}$ .
Down a group, A grows larger and the H–A bond weakens quickly, which outweighs the falling electronegativity: $\ce{HF} < \ce{HCl} < \ce{HBr} < \ce{HI}$ , and $\ce{H2O} < \ce{H2S} < \ce{H2Se}$ .

HF ionizes only partly in water (p $K_\mathrm{a}$ 3.17), partly because the H–F bond is very strong (565 kJ/mol) and $\ce{F-}$ hydrogen-bonds strongly to water. Its glass etching is another story: $\ce{SiO2} + \ce{6HF} \mathrel{\ce{->}} \ce{SiF6^2-} + \ce{2H3O+}$ is driven by the great stability of Si–F bonds, not by acid strength.

Lower pKa, stronger acid

Acid-strength ladder

Fully ionised in water (levelled)

No measurable acidity in water

\ce{HI}

≈−9.3

\ce{HBr}

≈−8.7

\ce{HCl}

≈−6.3

\ce{HF}

3.17

← strongerpKaweaker →

Why this order?

Down the group the H–X bond weakens sharply (bond enthalpies HF 565, HCl 431, HBr 366, HI 298 kJ/mol); this outweighs the falling electronegativity, so HI is strongest and HF weakest.

Step through the families; each shows “why this order” below. Hollow markers are estimates that cannot be measured directly in water.

Oxoacids, the inductive effect, and hydrated metal ions

One idea runs through many families: anything that pulls electron density out of the O–H bond, or stabilizes the anion’s charge, strengthens the acid.

In hypohalous acids HOX, a more electronegative X makes a stronger acid: HOCl (7.53) > HOBr (8.65) > HOI (10.64).
In carboxylic acids, a distant Cl or F withdraws electrons along the σ bonds (the inductive effect): acetic 4.76, chloroacetic 2.87, dichloroacetic 1.35, trichloroacetic 0.66.
In hydrated metal ions $\ce{[M(H2O)6]^{n+}}$ , a small, highly charged metal ion strongly polarizes the O–H bonds of its water ligands: $\ce{[Fe(H2O)6]^3+} \mathrel{\ce{<=>}} \ce{[Fe(H2O)5(OH)]^2+} + \ce{H+}$ has $\mathrm{p}K_\mathrm{a}$ about 2.2, like phosphoric acid, while $\ce{Na+}$ is barely acidic at all.

Worked example

pH of 0.10 M iron(III) nitrate

Treat $\ce{[Fe(H2O)6]^3+}$ as a monoprotic weak acid with $\mathrm{p}K_\mathrm{a} = 2.19$ and estimate the pH of a 0.10 M solution (ignore further hydrolysis).

In 1923 Lewis proposed a broader definition: an acid is an electron-pair acceptor and a base an electron-pair donor. $\ce{BF3} + \ce{NH3} \mathrel{\ce{->}} \ce{F3B-NH3}$ transfers no proton, yet it is a textbook acid–base reaction; metal ions binding ligands and $\ce{CO2}$ dissolving to form carbonic acid are Lewis acid–base chemistry too. The proton itself is the simplest Lewis acid.

Concept check

Which 0.1 M solution is most acidic?

\ce{NaNO3}

\ce{Mg(NO3)2}

\ce{Al(NO3)3}

\ce{KNO3}

Try the calculation

Hypochlorous acid has $\mathrm{p}K_\mathrm{a} = 7.53$ . What is its $K_\mathrm{a}$ ?

Key takeaways

Across a period rising electronegativity polarizes H–A and strengthens the acid; down a group weaker bonds make stronger acids.
More terminal oxygens and a more electronegative X make stronger oxoacids; inductive effects travel through σ bonds.
Small, highly charged metal ions make their aqua ions noticeably acidic (e.g. Fe³⁺, Al³⁺).
Acids stronger than H₃O⁺ are levelled in water; Lewis theory extends acidity to electron-pair acceptance.

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Data and references

D. C. Harris, Quantitative Chemical Analysis, 9th ed. (2016), Appendix G.
C. F. Baes, R. E. Mesmer, The Hydrolysis of Cations, Wiley (1976).
C. E. Housecroft, A. G. Sharpe, Inorganic Chemistry, 5th ed., Pearson (2018), Ch. 7 (acids, bases and ions in aqueous solution).
P. Atkins et al., Shriver & Atkins’ Inorganic Chemistry, 5th ed. (2010), Ch. 4.