Hydrogen Halides & Oxoacids of Halogens

What Are Hydrogen Halides?

Each halogen of Group 17 reacts directly with hydrogen to give a binary covalent hydride of the type $\ce{HX}$ — collectively the hydrogen halides $\ce{HF}$, $\ce{HCl}$, $\ce{HBr}$ and $\ce{HI}$. In the gaseous state these molecules are essentially covalent, but when dissolved in water they ionise to release $\ce{H+}$, and their aqueous solutions are the familiar hydrohalic acids (hydrofluoric, hydrochloric, hydrobromic and hydriodic acid).

All four halogens react with hydrogen, but the affinity for hydrogen decreases steadily from fluorine to iodine. That single fact — declining affinity down the group — is the seed from which every trend on this page grows: the strongest H–X bond ($\ce{HF}$) is also the most reluctant to release $\ce{H+}$, the most thermally stable, and the poorest reducing agent.

Preparation of Hydrogen Halides

A useful exam pattern: $\ce{HF}$ and $\ce{HCl}$ are made by displacing the volatile acid from its salt with a non-volatile, non-oxidising acid such as concentrated $\ce{H2SO4}$; but $\ce{HBr}$ and $\ce{HI}$ cannot be made this way because hot concentrated $\ce{H2SO4}$ oxidises the liberated $\ce{HBr}$/$\ce{HI}$ back to the free halogen. For the heavier two we use a non-oxidising acid ($\ce{H3PO4}$) or hydrolysis of the phosphorus trihalide.

Bond Enthalpy & Thermal Stability

As we descend the group, the H–X bond length increases ($\ce{H-F}$ 91.7 pm, $\ce{H-Cl}$ 127.4 pm, $\ce{H-Br}$ 141.4 pm, $\ce{H-I}$ 160.9 pm) because the halogen atom gets larger and its valence orbital overlaps the small hydrogen 1s orbital less effectively. A longer, weaker bond means the bond dissociation enthalpy falls in the order:

$$\ce{H-F > H-Cl > H-Br > H-I}$$

Thermal stability tracks the bond strength. Since $\ce{HF}$ holds its bond most tightly, it is the most stable to thermal decomposition, while $\ce{HI}$ decomposes most readily into its elements. The stability of the hydrogen halides therefore decreases in exactly the same order: $\ce{HF > HCl > HBr > HI}$.

Acid Strength: HF < HCl < HBr < HI

This is the single most-tested fact in the whole subtopic. The strength of a hydrohalic acid in water is governed chiefly by how easily the H–X bond breaks to release $\ce{H+}$ — that is, by the H–X bond dissociation enthalpy, not by the electronegativity of the halogen. As the halogen grows larger down the group, the H–X bond lengthens and weakens, the proton is lost more readily, and the acid strength rises:

$$\text{Acid strength: } \ce{HF < HCl < HBr < HI}$$

The $pK_a$ figures from the source make the contrast vivid: $\ce{HF}$ has a positive $pK_a$ (3.2) — it is only weakly ionised — whereas $\ce{HCl}$, $\ce{HBr}$ and $\ce{HI}$ all have strongly negative $pK_a$ values, with $\ce{HI}$ ($pK_a \approx -10$) the strongest of all. A smaller (more negative) $pK_a$ means a stronger acid.

Why HF Is the Odd One Out

Fluorine is the most electronegative element, so intuition wrongly suggests $\ce{HF}$ should be the strongest acid. In fact it is the weakest, and it breaks ranks on two fronts — both traceable to fluorine's tiny size and high electronegativity.

1. Weakest acid despite highest electronegativity

The $\ce{H-F}$ bond is the shortest and strongest of the four. Breaking it to liberate $\ce{H+}$ costs the most energy, so $\ce{HF}$ ionises only to a very small extent in water:

$$\ce{HF + H2O <=> H3O+ + F^-} \quad (\text{ionises only slightly})$$

The bond-enthalpy effect overrides electronegativity. This is why $\ce{HF}$, uniquely, is a weak acid while the other three are strong acids.

2. Abnormally high boiling point

$\ce{HF}$ is a liquid at room temperature (b.p. 293 K), while $\ce{HCl}$ (b.p. 189 K), $\ce{HBr}$ (206 K) and $\ce{HI}$ (238 K) are gases. The cause is intermolecular hydrogen bonding: the small, highly electronegative fluorine atom forms strong $\ce{F-H\bond{...}F}$ bridges, linking $\ce{HF}$ molecules into zig-zag chains that must be torn apart before the liquid can boil.

Reducing Character of HX

A reducing agent works by giving something up — here, by liberating hydrogen and being oxidised to the free halogen $\ce{X2}$. The easier the H–X bond breaks, the better the reductant. Since bond enthalpy falls HF→HI, the reducing power increases in the same direction:

$$\text{Reducing power: } \ce{HF < HCl < HBr < HI}$$

$\ce{HI}$, with its weak $\ce{H-I}$ bond, is the strongest reducing agent of the four and is used for that purpose in organic chemistry; $\ce{HF}$, with the strongest bond, is the poorest. Note that reducing power, thermal instability and acid strength all run the same way (HF→HI), while bond enthalpy and thermal stability run the opposite way.

Oxoacids of Halogens: The Map

When a halogen bonds to oxygen and an O–H group, the result is an oxoacid. In every halogen oxoacid the halogen is the central atom; chlorine, bromine and iodine each build a four-step ladder of oxoacids in oxidation states +1, +3, +5 and +7. Most cannot be isolated pure — they are stable only in aqueous solution or as their salts (hypochlorites, chlorites, chlorates, perchlorates and the analogous bromine/iodine salts).

Writing the chlorine oxoacids as $\ce{HOCl}$, $\ce{HOClO}$, $\ce{HOClO2}$ and $\ce{HOClO3}$ (rather than $\ce{HClO}$, $\ce{HClO2}$, $\ce{HClO3}$, $\ce{HClO4}$) is more honest about their structure: there is always exactly one O–H group, and the extra oxygens are terminal $\ce{Cl=O}$ atoms doubly bonded to chlorine. The number of terminal oxygens climbs 0 → 1 → 2 → 3 across the series.

Oxoacid Acid Strength & Oxidising Power

For a single halogen, the acid strength of its oxoacids increases with the number of oxygen atoms bonded to the central halogen — equivalently, with the oxidation state of the halogen. For chlorine:

$$\text{Acid strength: } \ce{HOCl < HOClO < HOClO2 < HOClO3}$$ $$\text{i.e. } \ce{HClO < HClO2 < HClO3 < HClO4}$$

The reasoning is electron withdrawal. Oxygen is more electronegative than chlorine, so each extra terminal oxygen pulls electron density away from the central chlorine, and hence away from the O–H bond. The O–H bond is progressively weakened, $\ce{H+}$ is released more easily, and the conjugate base (e.g. the $\ce{ClO4^-}$ ion) is increasingly stabilised by spreading the negative charge over more oxygens. $\ce{HClO4}$ needs the least energy to break its O–H bond and is therefore the strongest — indeed one of the strongest acids known — while $\ce{HOCl}$ is very weak.

Oxidising power runs the opposite way. The lower oxoacids — with the halogen in a low, easily-reduced oxidation state — are the better oxidisers: $\ce{HOCl}$ (and the hypochlorite ion) is a vigorous oxidising and bleaching agent because the $\ce{HOCl}$ formed in chlorine water supplies nascent oxygen, which is responsible for chlorine's bleaching action. Perchloric acid, with chlorine already at its maximum +7 state, is the strongest acid but a comparatively reluctant oxidiser in dilute solution.

Why Fluorine Has Only HOF

Fluorine is the lone exception to the four-rung oxoacid ladder. It forms only one oxoacid, hypofluorous acid $\ce{HOF}$ — also called fluoric(I) acid — in which fluorine is in the +1 state. Two reasons, both rooted in fluorine's nature, explain the restriction:

The contrast is the textbook explanation for why fluorine "forms only one oxoacid while other halogens form a number of oxoacids." Chlorine, bromine and iodine all have accessible d orbitals, can expand their octets, and so populate the full +1/+3/+5/+7 series.