Acid Strength Down a Group and Across a Period

By Aniket Bhardwaj · 18 September 2026 · Chemistry Concept

Ask a class why HI is a stronger acid than HF and someone will answer "because iodine is more electronegative". Iodine is less electronegative than fluorine, and HI is still the stronger acid by a huge margin. That single contradiction is the reason this topic produces so many wrong answers: the trend down a group and the trend across a period are driven by two different factors. Learn which factor rules where, and the whole chapter becomes predictable instead of confusing.

One idea unifies everything: conjugate base stability

HA ⇌ H⁺ + A⁻ · K_a = [H⁺][A⁻] ÷ [HA] · pK_a = −log K_a

A strong acid is one whose conjugate base A⁻ is stable and comfortable holding the negative charge, so the equilibrium sits to the right. A large K_a means a small pK_a — that inversion is worth saying aloud until it is automatic. Every explanation below reduces to the same question: how well does A⁻ cope with the extra electron?

Down a group — bond enthalpy wins

Go down group 17 and the H–X bond gets longer and much weaker, because the halogen's valence orbital is larger and overlaps the small hydrogen 1s orbital poorly. A weaker bond is easier to break, and the resulting large anion spreads its charge over a bigger volume, which is stabilising.

Acid	Mean H–X bond enthalpy (kJ/mol)	Electronegativity of X (Pauling)	Approximate pK_a in water	Strength
HF	567	4.0	≈ 3.2	Weak acid
HCl	431	3.0	≈ −7	Strong
HBr	366	2.8	≈ −9	Stronger
HI	298	2.5	≈ −10	Strongest

Read the table across and the point makes itself. Electronegativity falls down the group while acid strength rises — so electronegativity cannot be the controlling factor here. Bond enthalpy falls steeply, from 567 to 298 kJ/mol, and acid strength tracks it exactly.

An honest caution about those pK_a numbers. In water, HCl, HBr and HI are all completely ionised, so water cannot tell them apart — this is the levelling effect: no acid stronger than H₃O⁺ can survive in water. Their negative pK_a values are estimates obtained from measurements in less basic solvents (glacial acetic acid is the classic choice), so treat them as approximate ordering rather than precise data, and expect different textbooks to quote slightly different figures.

The same trend holds in group 16: H₂O < H₂S < H₂Se < H₂Te in acid strength, for exactly the same bond-strength reason. Water is a very weak acid; hydrogen telluride is a moderately strong one.

Across a period — electronegativity wins

Now move along period 2. The bonds are all to second-row atoms of similar size, so bond enthalpy no longer dominates. What changes sharply is electronegativity, and a more electronegative atom holds the negative charge of A⁻ much more comfortably.

Acid	Conjugate base	Electronegativity of the atom carrying the charge	Approximate pK_a
CH₄	CH₃⁻	C, 2.5	≈ 50 (tables vary widely)
NH₃	NH₂⁻	N, 3.0	≈ 38
H₂O	OH⁻	O, 3.5	15.7
HF	F⁻	F, 4.0	≈ 3.2

Methane is so weak an acid that its pK_a cannot be measured directly and is estimated by indirect routes — which is why different sources quote noticeably different numbers for it. The ordering, however, is not in doubt, and the ordering is what exams ask for.

The rule to write in an answer: across a period, acid strength is decided by the electronegativity of the atom bearing the negative charge. Down a group, it is decided by the H–X bond enthalpy and the size of the anion.

Oxyacids — two separate comparisons

For acids of the form H–O–X, the acidic hydrogen is always on oxygen, so the question becomes how strongly the rest of the molecule pulls electron density away from that O–H bond.

Same central atom, different number of oxygens: more oxygen atoms means more electron withdrawal and better spreading of the charge on the anion, so the acid is stronger.

Acid	Oxidation state of Cl	Terminal (non-OH) oxygens	Approximate pK_a
HOCl (hypochlorous)	+1	0	≈ 7.5
HClO₂ (chlorous)	+3	1	≈ 2.0
HClO₃ (chloric)	+5	2	≈ −1
HClO₄ (perchloric)	+7	3	≈ −8

Pauling noticed an empirical shortcut that reproduces this column surprisingly well:

pK_a ≈ 8 − 5q, where q = number of terminal oxygen atoms (oxygens not carrying an H)

Test it: HOCl has q = 0 → predicted 8, observed ≈ 7.5. HClO₂ has q = 1 → predicted 3, observed ≈ 2.0. HClO₃ has q = 2 → predicted −2, observed ≈ −1. HClO₄ has q = 3 → predicted −7, observed ≈ −8. It is a rule of thumb, not a law, and it should be labelled as such in an answer — but as an order-of-magnitude estimate it is excellent.

Different central atom, same number of oxygens: now electronegativity decides, so HOCl > HOBr > HOI in acid strength — chlorine pulls the hardest on the O–H bond.

Worked example 1 — pK_a to pH for a weak acid

Find the pH of 0.10 M acetic acid, pK_a = 4.76.

Step 1. K_a = 10^−4.76 = 10^0.24 × 10⁻⁵ = 1.74 × 10⁻⁵

Step 2. For a weak acid where dissociation is small, [H⁺] ≈ √(K_a × C)
[H⁺] = √(1.74 × 10⁻⁵ × 0.10) = √(1.74 × 10⁻⁶) = 1.32 × 10⁻³ M

Step 3. Check the approximation: 1.32 × 10⁻³ ÷ 0.10 = 1.3%, well under the usual 5% limit, so it is safe.

Step 4. pH = −log(1.32 × 10⁻³) = 3 − log 1.32 = 3 − 0.120 = 2.88

Worked example 2 — one chlorine atom, and the approximation breaks

Replace one hydrogen of the methyl group with chlorine and the electron-withdrawing effect travels along the σ bonds to stabilise the carboxylate anion. This is the inductive effect, and it is the organic version of everything said above.

Find the pH of 0.10 M chloroacetic acid, ClCH₂COOH, pK_a = 2.86.

Step 1. K_a = 10^−2.86 = 10^0.14 × 10⁻³ = 1.38 × 10⁻³ — about 79 times larger than acetic acid's, from a single chlorine atom.

Step 2 (approximate). [H⁺] ≈ √(1.38 × 10⁻³ × 0.10) = √(1.38 × 10⁻⁴) = 1.17 × 10⁻² M, giving pH = 1.93.

Step 3 — check it. 1.17 × 10⁻² ÷ 0.10 = 11.7%, which is well above 5%. The approximation is not valid here, so solve properly:

x² ÷ (0.10 − x) = 1.38 × 10⁻³ → x² + 1.38 × 10⁻³x − 1.38 × 10⁻⁴ = 0
Discriminant = (1.38 × 10⁻³)² + 4(1.38 × 10⁻⁴) = 1.91 × 10⁻⁶ + 5.52 × 10⁻⁴ = 5.541 × 10⁻⁴
√(5.541 × 10⁻⁴) = 2.354 × 10⁻²
x = (2.354 × 10⁻² − 1.38 × 10⁻³) ÷ 2 = 2.216 × 10⁻² ÷ 2 = 1.108 × 10⁻² M

pH = −log(1.108 × 10⁻²) = 2 − 0.045 = 1.96

The shortcut gave 1.93 and the correct method gives 1.96. Small here — but the gap grows fast for stronger weak acids, and examiners choose exactly those numbers.

The inductive series, and why HF breaks the halide pattern

Adding more electron-withdrawing groups keeps pushing acidity up. Commonly tabulated pK_a values at 25 °C: CH₃COOH 4.76, ClCH₂COOH 2.86, Cl₂CHCOOH 1.29, Cl₃CCOOH 0.66. Each chlorine drags more electron density away from the carboxylate, spreading the charge and stabilising the conjugate base.

Finally, the anomaly students are most often asked to explain: why is HF a weak acid when fluorine is the most electronegative element? Two reasons combine. The H–F bond is by far the strongest of the four at 567 kJ/mol, so breaking it is expensive. And F⁻ is very small and highly charge-dense, so in water it forms unusually strong hydrogen bonds and ion pairs with H₃O⁺, pulling the equilibrium back towards undissociated HF. Both effects work against ionisation, and they outweigh the electronegativity advantage.

Common mistakes that cost marks

Using electronegativity to explain a group trend. It predicts the opposite of reality for HF → HI. Down a group, argue from bond enthalpy and anion size.
Using bond enthalpy to explain a period trend. Across period 2 the bonds are too similar; electronegativity of the charge-carrying atom is the argument.
Confusing strength with concentration. A concentrated solution of a weak acid can have a lower pH than a dilute strong acid. Strength is K_a; concentration is molarity.
Reading pK_a the wrong way round. Lower pK_a = stronger acid, always.
Skipping the 5% check before using [H⁺] = √(K_aC). If dissociation exceeds about 5%, the quadratic is compulsory.
Forgetting the levelling effect and claiming you can measure the difference between HCl and HI in water. In water they are equally strong; the difference shows only in a weaker base solvent.

Where this appears in exams

Exam	Typical question
CBSE / ICSE Class 11	Arrange HF, HCl, HBr, HI in order of acid strength and justify
CBSE / ICSE Class 12	Acidity of phenols vs alcohols vs carboxylic acids; effect of substituents
JEE / NEET	Ordering oxyacids; pH of weak acid solutions; inductive effect ranking
IIT-JAM / CUET-PG	K_a and pK_a calculations, conjugate acid–base pairs
GATE / CSIR-NET	Hammett-type substituent effects, acidity in non-aqueous solvents

Check your pH working in seconds. The pH / pOH calculator takes [H⁺], [OH⁻], pH or pOH and converts between all four, so you can verify both worked examples above — including the one where the square-root shortcut fails and the quadratic is needed.

Open the pH / pOH Calculator →

Periodic trends reward understanding over memorising, and that is much easier with a teacher who asks you "why" every time. ABC Chemistry runs Class 11–12 chemistry coaching at the Gurugram centre and online classes across India — abcchemistry.in.