The Strong Acids (and Why There Are So Few)

There are millions of acids. Every carboxylic acid, every phenol, every hydrated metal ion and thousands of inorganic compounds can donate a proton. Yet when textbooks list the strong acids, the list usually has just six or seven entries. That short list is worth memorising, because once you know it, you know that almost every other acid you meet is weak.

What “strong” means

A strong acid ionises completely in water. In a dilute solution, essentially every molecule has given its proton to a water molecule:

HA + H₂O → H₃O⁺ + A⁻

There’s a single arrow, not an equilibrium arrow, because there’s practically no HA left. In contrast, a weak acid sets up an equilibrium in which most molecules stay intact. If you’re unsure of the distinction, read concentrated vs strong first, because “strong” has nothing to do with how much acid is in the bottle.

The list

Acid	Formula	Estimated pKa	Conjugate base
Hydrochloric acid	HCl	about −6	Cl⁻ (chloride)
Hydrobromic acid	HBr	about −9	Br⁻ (bromide)
Hydroiodic acid	HI	about −10	I⁻ (iodide)
Nitric acid	HNO₃	about −1.4	NO₃⁻ (nitrate)
Sulfuric acid (first proton)	H₂SO₄	about −3	HSO₄⁻ (hydrogensulfate)
Perchloric acid	HClO₄	about −10 or lower	ClO₄⁻ (perchlorate)
Chloric acid	HClO₃	about −1	ClO₃⁻ (chlorate)

Negative pKa values can’t be measured directly in water (more on that below), so these figures are estimates from other solvents and theoretical calculations. Different sources quote slightly different numbers. What matters is that they’re all well below zero.

Some courses stop at six and leave out chloric acid. Some add less common strong acids such as permanganic acid (HMnO₄). If your syllabus gives a list, learn that one.

A closer look at each acid

Hydrochloric acid, HCl

A solution of hydrogen chloride gas in water. It’s the acid in your stomach, at roughly 0.01–0.1 mol/dm³. Industrially it’s used to “pickle” steel (dissolve rust and scale from its surface), to make PVC precursors, and to adjust pH in water treatment. Concentrated hydrochloric acid is about 37% HCl by mass (around 12 mol/dm³) and gives off choking fumes.

Hydrobromic and hydroiodic acids, HBr and HI

Chemically similar to hydrochloric acid but even stronger, for reasons explained below. They’re used mainly in organic synthesis, for example to convert alcohols into bromoalkanes and iodoalkanes. Hydroiodic acid is also a strong reducing agent and slowly turns brown in air as iodide is oxidised to iodine.

Nitric acid, HNO₃

Colourless when pure but often yellowish from dissolved nitrogen dioxide. Besides being a strong acid, it’s a strong oxidising agent, so it dissolves copper and silver, which ordinary acids can’t. It’s made by the Ostwald process from ammonia and is used mainly to make ammonium nitrate fertiliser.

Sulfuric acid, H₂SO₄

The most-produced chemical in the world by mass. Its first proton is fully released; the second (HSO₄⁻ → SO₄²⁻) is only partly released, with pKa2 of about 2. So sulfuric acid is a strong acid, but it isn’t “doubly strong”. The details are in polyprotic acids. Concentrated sulfuric acid is also a powerful dehydrating agent: it chars sugar to black carbon by pulling out the elements of water.

Perchloric acid, HClO₄

One of the strongest ordinary acids known. It’s used in analytical chemistry and to make perchlorate salts, including ammonium perchlorate for solid rocket fuel. Hot, concentrated perchloric acid is a dangerous oxidiser and needs special fume cupboards, because perchlorate residues can explode.

Chloric acid, HClO₃

Exists only in solution and is unstable when concentrated. It’s rarely used directly, but its salts, the chlorates, are important oxidisers.

Why hydrofluoric acid isn’t strong

The halogen acids HCl, HBr and HI are all strong, so it seems natural that HF would be too. It isn’t. HF has a pKa of about 3.2, making it a weak acid: in a 0.1 mol/dm³ solution, fewer than one molecule in ten is ionised. Two reasons:

1. The H–F bond is very strong. Fluorine is small, and its bond to hydrogen is short and strong (about 570 kJ/mol, compared with 432 for H–Cl). It takes more energy to break.
2. F⁻ is very tightly solvated and forms strong hydrogen bonds. In solution, the fluoride ion and undissociated HF bond strongly to each other and to water, holding the proton back.

Down the group, bonds get longer and weaker (H–Br about 366 kJ/mol, H–I about 298 kJ/mol), so acid strength increases: HF < HCl < HBr < HI.

Weak doesn’t mean safe. Hydrofluoric acid is one of the most dangerous acids in any lab. The undissociated HF molecule passes through skin easily and releases fluoride ions deep in the tissue, where they bind calcium. Burns may not hurt at first but can cause serious harm. It’s never used in schools.

Why oxyacids become strong

For acids with oxygen, strength depends on how many oxygen atoms are attached to the central atom without hydrogen on them. Each extra oxygen pulls electron density away from the O–H bond and spreads the negative charge of the conjugate base over more atoms. That’s why the chlorine oxyacids get dramatically stronger as oxygens are added:

Acid	Formula	pKa
Hypochlorous	HOCl	7.5 (weak)
Chlorous	HClO₂	about 2 (weak)
Chloric	HClO₃	about −1 (strong)
Perchloric	HClO₄	about −10 (very strong)

A useful rule of thumb: in an oxyacid written as (HO)ₘXOₙ, the acid is strong if n is 2 or more. Nitric acid (HO)NO₂ has n = 2; sulfuric acid (HO)₂SO₂ has n = 2; perchloric acid (HO)ClO₃ has n = 3. Phosphoric acid (HO)₃PO has n = 1, so it’s weak.

The levelling effect: why they all look equally strong

In water, you can’t tell which strong acid is strongest. Hydrochloric, nitric and perchloric acids at 0.1 mol/dm³ all have pH 1.0.

That’s because water is a base, and every one of these acids hands its proton over completely. The strongest acid that can exist in water is H₃O⁺ itself. Anything stronger simply converts to H₃O⁺. Chemists call this the levelling effect: water levels all strong acids down to the strength of H₃O⁺.

To rank them, you have to use a solvent that’s a weaker base than water, such as pure ethanoic acid. In that solvent the acids don’t all ionise completely, and differences appear: perchloric acid is stronger than sulfuric acid, which is stronger than hydrochloric acid, which is stronger than nitric acid.

Stronger still: superacids

Acids stronger than 100% sulfuric acid are called superacids. Examples include triflic acid (CF₃SO₃H), fluorosulfuric acid, and mixtures such as “magic acid” (fluorosulfuric acid with antimony pentafluoride). They can protonate substances normally considered non-basic, including hydrocarbons. George Olah won the 1994 Nobel Prize in Chemistry partly for using superacids to study carbocations.

Memory tips

A popular way to remember the classic six: “So I Brought No Clean Clothes”, standing for Sulfuric, hydroIodic, hydroBromic, Nitric, hydroChloric and perChloric. Any mnemonic you invent yourself will stick better, though.

Another shortcut: if an acid isn’t on your strong list, treat it as weak. That will be right almost every time at school level.

Key takeaways

• Strong acids ionise completely in water. The common ones are HCl, HBr, HI, HNO₃, H₂SO₄ (first proton), HClO₄ and HClO₃.
• HF is weak because of its very strong H–F bond and strong hydrogen bonding, but it’s extremely hazardous.
• Oxyacids get stronger as more non-hydroxyl oxygens are attached to the central atom.
• Water levels all strong acids to the strength of H₃O⁺; other solvents reveal their differences.
• Calculate their pH directly from concentration with the pH calculator, and compare them with weak acids in strong vs weak acids.