The p-Block Elements: Halogens and Noble Gases

Group 17, the halogens, share the valence configuration ns2 np5, exactly one electron short of a complete, stable octet, and this single missing electron is the root cause of nearly every characteristic halogen property: high electronegativity, strong oxidising power, and a strong drive to gain that one electron and become a stable, singly-charged halide ion. Oxidising power falls steadily down the group, F2 > Cl2 > Br2 > I2, tracking each element's shrinking pull on an incoming electron as atomic size grows and the extra electron ends up further from the nucleus. This trend is genuinely useful predictively: any halogen higher in the group can displace a halide ion lower in the group from solution (chlorine displacing bromide or iodide, for instance), since the higher halogen is the stronger oxidising agent of the two.

Fluorine breaks the smooth trend it otherwise sets: the F-F bond is actually weaker than the Cl-Cl bond, a genuinely counterintuitive result given how small and how strongly electronegative fluorine is. The explanation lies in fluorine's very small size: its lone pairs sit unusually close to the bonding pair, and the resulting lone-pair-to-lone-pair repulsion between the two fluorine atoms significantly weakens the F-F bond compared to what you would otherwise predict. This weak bond is precisely why fluorine is so ferociously reactive despite having a comparatively low bond dissociation energy to overcome, an important distinction between kinetic reactivity (how fast a reaction actually proceeds, partly governed by how easily the starting bond breaks) and simple thermodynamic bond strength; fluorine's weak F-F bond breaks easily, so fluorine gas reacts explosively with almost everything it meets, including glass and even some noble gases.

The acid strength of the hydrogen halides increases down the group, HF < HCl < HBr < HI, which surprises many students at first, since fluorine is by far the most electronegative halogen and might seem like it should make HF the strongest acid. The deciding factor is instead H-X bond dissociation energy, which falls steadily down the group as atomic size grows and the bonding electron pair sits further from each nucleus (H-F is roughly 570 kJ/mol, while H-I is only about 299 kJ/mol), and a weaker bond ionises far more readily in water, releasing H+ more easily regardless of how electronegative the halogen itself is. HF is genuinely the anomaly among the four: it is actually a weak acid in water, and the reason is its own extensive hydrogen bonding, both within liquid HF itself and with surrounding water molecules, which stabilises the molecular H-F species and resists its own ionisation.

Every halogen shows the -1 oxidation state, gaining one electron to complete its octet, but chlorine, bromine, and iodine additionally show a range of positive oxidation states, +1, +3, +5, and +7, most visibly in their oxoacids and oxoanions, made possible by their accessible valence-shell d orbitals, which let them promote and unpair electrons to form more than one covalent bond to oxygen. Fluorine is the sole exception: it never shows a positive oxidation state at all, for two compounding reasons, it has no d orbitals available in its valence shell (n=2 has only s and p) to promote electrons into, and it is simultaneously the most electronegative element in the entire periodic table, so it can never be forced to share electrons unequally in oxygen's favour. This is exactly the same underlying limitation already seen with boron's capped coordination number and nitrogen's own capped covalency, a genuinely recurring theme across the whole of this syllabus: second-period elements lacking valence d orbitals behave in ways their heavier group-mates simply cannot.

Chlorine forms a genuinely instructive series of four oxoacids across its accessible positive oxidation states: hypochlorous acid (HOCl, +1), chlorous acid (HOClO, +3), chloric acid (HOClO2, +5), and perchloric acid (HOClO3, +7), and acid strength climbs steadily across the series in exactly that order, HOCl being the weakest and HClO4 among the strongest acids known. The explanation is the same resonance-stabilisation logic covered for other oxoacid families: each additional oxygen atom (beyond the one already bonded to the acidic hydrogen) is a strongly electron-withdrawing group that pulls electron density away from the O-H bond, weakening it and making the proton easier to lose, while also spreading the resulting negative charge across more resonance structures once that proton is lost, further stabilising the conjugate base. Perchloric acid's own conjugate base, perchlorate (ClO4-), is an exceptionally stable, symmetrical tetrahedral ion with four genuinely equivalent Cl-O bonds by resonance, which is exactly why perchloric acid ranks among the strongest acids known.

Interhalogen compounds form between two different halogens, and they fall into four structural families based on the ratio of atoms involved, AB (such as ClF), AB3 (such as ClF3), AB5 (such as ClF5 or BrF5), and AB7 (the single example, IF7), with the larger halogen always occupying the central position and the smaller, more electronegative halogen(s) surrounding it, exactly as VSEPR theory would predict from central-atom size and available d orbitals. Every interhalogen compound is, without exception, more reactive than the pure halogen element made only of the larger, central atom, and the reason is that the A-B bond in an interhalogen is inherently weaker and more polar than the A-A bond in the pure halogen, since the two different-sized, different-electronegativity atoms fit together and share electrons less efficiently than two identical atoms do, making that bond correspondingly easier to break in a subsequent reaction.

Bleaching powder, calcium oxychloride (CaOCl2), is prepared industrially by passing chlorine gas over dry slaked lime, Ca(OH)2 + Cl2 to CaOCl2 + H2O, and its bleaching action, along with its usefulness as a disinfectant, comes from the same underlying chemistry: on exposure to atmospheric carbon dioxide or dilute acid, bleaching powder slowly liberates chlorine, which reacts further with moisture in the air to release nascent (freshly formed, unusually reactive) oxygen, and it is this nascent oxygen that actually oxidises coloured organic compounds to colourless ones, and destroys the cell walls of microorganisms, giving bleaching powder its practical value in both textile bleaching and water disinfection alike.

Group 18, the noble gases (helium, neon, argon, krypton, xenon, and radon), share a complete valence configuration, ns2 np6 for every member except helium (1s2), and this fully satisfied octet is exactly why they were, for decades after their discovery by William Ramsay and collaborators in the 1890s, assumed to be totally chemically inert, giving the whole group its original name, the 'inert gases'. Their everyday uses trace directly back to this same unreactivity and to characteristic physical properties: helium's very low density and total non-flammability make it ideal for balloons and airships (unlike the far more dangerous, flammable hydrogen once used for the same purpose) and, mixed with oxygen, for deep-sea diving gas mixtures (avoiding the nitrogen narcosis and decompression sickness risks of ordinary compressed air); neon's characteristic reddish-orange glow under electrical discharge gives it the classic neon-sign application; and argon's total inertness makes it the standard shielding gas for arc welding and for filling incandescent light bulbs, protecting the hot metal filament or weld from unwanted reaction with atmospheric oxygen.

In 1962, Neil Bartlett prepared the first genuine noble gas compound, xenon hexafluoroplatinate, overturning decades of assumed total inertness at a stroke, and xenon remains, together with the far rarer and radioactive radon, essentially the only noble gas that forms a substantial, well-characterised chemistry, reacting directly only with the two most electronegative elements, fluorine and oxygen. The reason xenon specifically, rather than helium, neon, argon, or krypton, is the answer lies in a straightforward trend: ionisation enthalpy falls steadily down group 18 as atomic size increases, and xenon's outermost electrons are held loosely enough (its first ionisation enthalpy is lower than that of oxygen molecules, remarkably) that a sufficiently powerful oxidiser like fluorine can genuinely pull one away and force xenon into forming real covalent bonds, an option that simply is not available to helium, neon, or argon, whose electrons remain held far too tightly.

Xenon reacts directly with fluorine gas under controlled conditions to form three distinct binary fluorides, XeF2, XeF4, and XeF6, and each one's shape follows directly and predictably from VSEPR theory, provided xenon's own lone pairs (it starts with four, in its free atomic state) are counted alongside its bonding pairs. XeF2 has two Xe-F bonding pairs and three lone pairs, five electron domains total, giving a trigonal bipyramidal electron geometry with the three lone pairs occupying the roomier equatorial positions and the two fluorine atoms left in the axial positions, producing an overall linear molecular shape. XeF4 has four Xe-F bonding pairs and two lone pairs, six electron domains total, giving an octahedral electron geometry with the two lone pairs positioned opposite each other (minimising their mutual repulsion) and the four fluorine atoms in the remaining positions, producing a square planar molecular shape. XeF6 has six Xe-F bonding pairs and one remaining lone pair, seven electron domains total, giving a shape usually described as a distorted octahedron, since one lone pair among six identical bonding pairs pushes the surrounding fluorine atoms noticeably out of a perfect octahedral arrangement.