Group 17 — Halogen Family

Occurrence and electronic configuration

The name halogen comes from the Greek halo (salt) and genes (born) — the "salt-producers". The group comprises fluorine, chlorine, bromine, iodine, astatine and the synthetic tennessine; astatine and tennessine are radioactive, and tennessine's chemistry has never been established because only milligram-scale amounts with millisecond half-lives exist. Members of Group 17 show closer family resemblance, with a more regular gradation of properties, than any other group of the periodic table.

Fluorine and chlorine are fairly abundant; bromine and iodine less so. Fluorine occurs mainly as insoluble fluorides — fluorspar $\ce{CaF2}$, cryolite $\ce{Na3AlF6}$ and fluoroapatite $\ce{3Ca3(PO4)2.CaF2}$. Sea water is essentially a sodium chloride solution (about 2.5% by mass) but also carries bromides and iodides, while dried-up sea deposits supply carnallite $\ce{KCl.MgCl2.6H2O}$. Iodine appears in marine life: certain seaweeds contain up to 0.5% iodine and Chile saltpetre up to 0.2% sodium iodate.

Every halogen has the valence configuration $ns^2np^5$ — seven electrons in the outermost shell, exactly one short of the next noble gas. That single-electron deficit is the master key to almost all halogen chemistry: it explains the high electronegativity, the strongly negative electron gain enthalpy and the dominant $-1$ oxidation state.

Atomic radii, ionisation enthalpy and electronegativity

Halogens have the smallest atomic radii in their respective periods because the maximum effective nuclear charge for that period pulls the valence shell inward. Fluorine, being a second-period element, is extremely small. Both atomic and ionic radii increase from F to I as successive quantum shells are added.

Because each atom is only one electron short of a stable octet, halogens have very little tendency to lose electrons, hence very high ionisation enthalpy. Ionisation enthalpy falls down the group as the atom enlarges and the outer electrons become easier to remove. Electronegativity follows the same downward trend: fluorine is the most electronegative element in the periodic table.

*Fluorine is conventionally assigned 4.0 on the Pauling scale, the highest of any element; the supplementary table tabulates Cl 3.2, Br 3.0, I 2.7. All other figures are from Table 7.8 of the supplementary text.

Electron gain enthalpy and the Cl > F anomaly

Halogens have the most negative (maximum) electron gain enthalpy in their respective periods, because adding one electron completes a stable noble-gas configuration. Down the group the electron gain enthalpy becomes less negative as size increases. But there is a celebrated exception at the top: the negative electron gain enthalpy of fluorine is smaller than that of chlorine ($-333$ vs $-349$ kJ mol⁻¹).

The reason is fluorine's tiny size. Its 2p orbitals are so compact that the electrons already present crowd together, producing strong inter-electronic repulsions. An incoming electron entering these small 2p orbitals is repelled and does not feel the full nuclear attraction, so less energy is released than in chlorine, whose larger 3p orbitals accommodate the extra electron more comfortably. Chlorine therefore holds the record for the most negative electron gain enthalpy of all the elements.

Bond enthalpy and the F₂ anomaly

A second curiosity sits in the X–X bond enthalpies. One would expect bond strength to fall steadily down the group, yet the dissociation enthalpy of $\ce{F2}$ (158.8 kJ mol⁻¹) is smaller than that of $\ce{Cl2}$ (242.6) and $\ce{Br2}$ (192.8). From chlorine onward the expected order returns: $\text{Cl–Cl} > \text{Br–Br} > \text{I–I}$. The complete order is therefore

$\text{Cl–Cl} > \text{Br–Br} > \text{F–F} > \text{I–I}$

The anomaly arises because the two fluorine atoms in $\ce{F2}$ are so small that their non-bonding lone pairs are forced very close together. The resulting large lone-pair / lone-pair (electron-electron) repulsion weakens the F–F bond — an effect far less pronounced in the larger $\ce{Cl2}$ molecule.

Physical states, colour and oxidation states

Halogens display a smooth gradation of physical properties. Down the group fluorine and chlorine are gases, bromine a liquid and iodine a solid, with melting and boiling points rising steadily as molar mass and van der Waals forces increase. All halogens are coloured: they absorb radiation in the visible region, exciting outer electrons to higher levels, and the quantum of radiation absorbed sets the colour seen — yellow $\ce{F2}$, greenish-yellow $\ce{Cl2}$, red $\ce{Br2}$ and violet $\ce{I2}$. Bromine and iodine are only sparingly soluble in water but dissolve readily in organic solvents such as chloroform and carbon tetrachloride to give coloured solutions.

All halogens exhibit the $-1$ oxidation state. Chlorine, bromine and iodine additionally show $+1, +3, +5$ and $+7$ states because their atoms can be promoted to excited states with three, five and seven unpaired electrons, the extra electrons occupying available $d$ orbitals. These higher states appear only in combination with the small, highly electronegative oxygen and fluorine atoms — in interhalogens, oxides and oxoacids. Fluorine, lacking $d$ orbitals and being the most electronegative element, can neither expand its octet nor take a positive oxidation state, so it is restricted to $-1$.

Oxidising power and reactivity

The ready acceptance of an electron makes all halogens strong oxidising agents, and reactivity decreases down the group. $\ce{F2}$ is the strongest oxidising halogen: a halogen will oxidise the halide ion of any halogen of higher atomic number.

$\ce{F2 + 2X^- -> 2F^- + X2}$ (X = Cl, Br or I)
$\ce{Cl2 + 2X^- -> 2Cl^- + X2}$ (X = Br or I)
$\ce{Br2 + 2I^- -> 2Br^- + I2}$

The decreasing oxidising ability in aqueous solution is mirrored by the standard electrode potentials $E^\circ(\ce{X2}/\ce{X^-})$, which fall from $+2.87$ V for fluorine to $+0.54$ V for iodine. The electrode potential depends on three energy terms in sequence: bond dissociation enthalpy, electron gain enthalpy and hydration enthalpy of the halide ion.

The same trend shows in reactions with water. Fluorine oxidises water all the way to oxygen, whereas chlorine and bromine only partly react to give a hydrohalic and a hypohalous acid; iodine cannot oxidise water at all — in fact dissolved oxygen oxidises $\ce{I^-}$ in acid solution, the exact reverse of the fluorine reaction.

$\ce{2F2 + 2H2O -> 4H^+ + 4F^- + O2}$
$\ce{X2 + H2O -> HX + HOX}$ (X = Cl or Br)
$\ce{4I^- + 4H^+ + O2 -> 2I2 + 2H2O}$

Anomalous behaviour of fluorine

As the head element of the group and a second-period p-block element, fluorine breaks the smooth trends in several ways. Its ionisation enthalpy, electronegativity and electrode potential are all higher than the group trend would predict, while its ionic and covalent radii, melting and boiling points, bond dissociation enthalpy and electron gain enthalpy are all lower than expected. The root causes are its small size, highest electronegativity, low F–F bond dissociation enthalpy and the absence of $d$ orbitals in its valence shell.

Hydrogen fluoride is a liquid (b.p. 293 K) because the small, highly electronegative fluorine forms strong intermolecular hydrogen bonds; the larger, less electronegative halogens give gaseous $\ce{HCl}$, $\ce{HBr}$ and $\ce{HI}$. Note also the acid-strength order $\ce{HF} < \ce{HCl} < \ce{HBr} < \ce{HI}$, which tracks the falling H–X bond dissociation enthalpy down the group.

Chlorine — preparation, properties and uses

Chlorine was discovered by Scheele in 1774 from the action of $\ce{HCl}$ on $\ce{MnO2}$; Davy established its elementary nature in 1810 and named it after its colour (Greek chloros, greenish-yellow).

Preparation and manufacture

In the laboratory chlorine is made by oxidising concentrated $\ce{HCl}$ (or a salt + $\ce{H2SO4}$ mixture) with $\ce{MnO2}$ or $\ce{KMnO4}$:

$\ce{MnO2 + 4HCl -> MnCl2 + Cl2 + 2H2O}$
$\ce{4NaCl + MnO2 + 4H2SO4 -> MnCl2 + 4NaHSO4 + 2H2O + Cl2}$
$\ce{2KMnO4 + 16HCl -> 2KCl + 2MnCl2 + 8H2O + 5Cl2}$

Industrially it is obtained by Deacon's process — air oxidation of $\ce{HCl}$ over a $\ce{CuCl2}$ catalyst at 723 K — and chiefly by the electrolysis of brine (concentrated $\ce{NaCl}$ solution), where chlorine is liberated at the anode. It also arises as a by-product in many chemical industries.

$\ce{4HCl + O2 ->[CuCl2][723\,K] 2Cl2 + 2H2O}$

Properties

Chlorine is a greenish-yellow gas with a pungent, suffocating odour, about 2.5 times heavier than air, easily liquefied to a greenish-yellow liquid (b.p. 239 K) and soluble in water. It reacts directly with metals and non-metals to give chlorides, and has a strong affinity for hydrogen:

$\ce{2Al + 3Cl2 -> 2AlCl3}$    $\ce{P4 + 6Cl2 -> 4PCl3}$
$\ce{2Fe + 3Cl2 -> 2FeCl3}$    $\ce{H2 + Cl2 -> 2HCl}$

With alkalis the outcome depends on conditions — dilute/cold alkali gives chloride + hypochlorite, hot/concentrated alkali gives chloride + chlorate. Both are disproportionations, with chlorine going from the 0 state to $-1$ and $+1$ (or $+5$):

$\ce{2NaOH + Cl2 -> NaCl + NaOCl + H2O}$ (cold, dilute)
$\ce{6NaOH + 3Cl2 -> 5NaCl + NaClO3 + 3H2O}$ (hot, conc.)

With dry slaked lime it gives bleaching powder, $\ce{Ca(OCl)2.CaCl2.Ca(OH)2.2H2O}$:

$\ce{2Ca(OH)2 + 2Cl2 -> Ca(OCl)2 + CaCl2 + 2H2O}$

Oxidising and bleaching action — nascent oxygen

Chlorine water on standing loses its yellow colour as it forms $\ce{HCl}$ and hypochlorous acid $\ce{HOCl}$; the $\ce{HOCl}$ supplies nascent oxygen $[\ce{O}]$, which is responsible for both the oxidising and bleaching properties of chlorine. As an oxidant, chlorine converts ferrous to ferric, sulphite to sulphate, $\ce{SO2}$ to $\ce{H2SO4}$ and $\ce{I2}$ to iodate:

$\ce{2FeSO4 + H2SO4 + Cl2 -> Fe2(SO4)3 + 2HCl}$
$\ce{SO2 + 2H2O + Cl2 -> H2SO4 + 2HCl}$
$\ce{I2 + 6H2O + 5Cl2 -> 2HIO3 + 10HCl}$

Its bleaching action is due to oxidation and requires moisture. Nascent oxygen oxidises the coloured matter to a colourless product, and because the colour is destroyed rather than masked, the bleaching is permanent:

$\ce{Cl2 + H2O -> 2HCl + [O]}$
Coloured substance $+\ [\ce{O}] \longrightarrow$ Colourless substance

Uses

Chlorine is used for bleaching wood-pulp (paper and rayon) and cotton textiles; in the extraction of gold and platinum; in manufacturing dyes, drugs and organic compounds such as $\ce{CCl4}$, $\ce{CHCl3}$, DDT and refrigerants; in sterilising drinking water; and in preparing poisonous gases including phosgene $\ce{COCl2}$, tear gas $\ce{CCl3NO2}$ and mustard gas $\ce{ClCH2CH2SCH2CH2Cl}$.