Group 17 Elements: The Halogens -- General Trends

Group 17, the halogens, comprises fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and

the rare, intensely radioactive astatine (At), sharing the general outer configuration

$ns^2np^5$ -- just one electron short of a complete octet. This single structural fact

defines the entire group's chemical personality: every halogen atom has an exceptionally strong

"pull" toward gaining one more electron to complete its octet, making the halogens, as a group,

the most reactive nonmetals and the strongest common family of oxidising agents in the periodic

table.

Down the group, atomic and ionic radius increases steadily as successive shells are added,

and ionisation enthalpy decreases correspondingly, from fluorine's very high value down to

iodine's comparatively modest one. Electronegativity is at its overall maximum at fluorine

(the single most electronegative element in the entire periodic table) and falls steadily down

the group to iodine. Oxidising power falls in the same order, $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$: each halogen is able to oxidise (displace) the halide ion of any

halogen below it from solution (e.g. chlorine water displaces bromine from bromide solution),

but not the reverse.

Two closely related trends show a well-known anomaly at fluorine, both tracing back to

fluorine's unusually small atomic size. The electron gain enthalpy (the energy released when

a gaseous atom gains one electron) is, counter-intuitively, less negative for fluorine than for

chlorine (fluorine about $-328\ \text{kJ mol}^{-1}$ versus chlorine's more negative

$-349\ \text{kJ mol}^{-1}$), even though electronegativity itself is highest for fluorine. The

reason is that fluorine's valence shell ($n=2$) is so compact that the incoming extra electron

experiences unusually strong repulsion from the electrons already present, partly offsetting the

favourable pull of the nucleus; chlorine, with a larger $n=3$ valence shell and correspondingly

less inter-electron crowding, can accommodate the extra electron more favourably even though its

nuclear attraction is, in isolation, weaker than fluorine's. The bond dissociation enthalpy

of $\text{F}_2$ is similarly anomalously low (about $155\ \text{kJ mol}^{-1}$, distinctly lower

than $\text{Cl}_2$'s $242\ \text{kJ mol}^{-1}$) for essentially the same structural reason: the

two fluorine atoms in $\text{F}_2$ are so close together that their several non-bonding lone

pairs experience significant mutual repulsion across the short F-F bond, weakening the bond

relative to what its small size alone would predict; in the larger $\text{Cl}_2$ molecule, this

lone-pair repulsion is much less severe because the atoms (and therefore their lone pairs) sit

farther apart. This is also, notably, the reason fluorine is more reactive than its low bond

dissociation enthalpy alone would suggest is unusual -- the weak F-F bond is easily broken, …