p-Block Elements: the Oxygen and Halogen Groups and the Noble Gases

This chapter completes the systematic survey of the p-block main-group elements, turning to groups 16, 17 and 18. The same two organising mechanisms from the previous chapter — the first-row anomaly and the inert pair effect — continue to operate, but group 17 adds a third, decisive theme: fluorine's unique combination of small size and extreme electronegativity makes it behave unlike any of its own heavier congeners, not merely "the lightest halogen." Group 18 closes the block with the most dramatic overturning of textbook dogma in all of inorganic chemistry: the "inert" gases are not, in fact, universally inert.

You should have to hand: the first-row anomaly and inert pair effect from the previous chapter, VSEPR theory including lone-pair-containing geometries, and the general p-block oxide acidity trend.

The organising principle of this chapter. Three ideas recur: (1) the first-row anomaly, now showing up as oxygen's coordination-number ceiling and fluorine's complete inability to reach a positive oxidation state; (2) electronegativity and bond-strength anomalies specific to the lightest member of groups 16 and 17 (weak O–O and F–F single bonds from lone-pair repulsion, despite their position at the top of the group); (3) the systematic, predictable reactivity of the noble gases, governed by the same ionization-energy logic used throughout this course, once the historical assumption of total inertness is set aside.

1. Group 16: Oxidation States and the Oxygen/Sulfur Contrast

Group 16 ($ns^2np^4$) spans oxygen and sulfur (nonmetals) through selenium and tellurium (metalloids) to polonium (a radioactive metal). The group oxidation state is $+6$ (e.g. $\mathrm{SF_6}$, $\mathrm{H_2SO_4}$), but oxygen is essentially restricted to $-2$, $-1$ (peroxides), and $-\tfrac{1}{2}$ (superoxides) in its compounds, together with $0$ in $\mathrm{O_2}$/$\mathrm{O_3}$: oxygen shows a positive oxidation state only in fluorides (e.g. $\mathrm{OF_2}$, O is $+2$; $\mathrm{O_2F_2}$, O is $+1$) and in the dioxygenyl cation $\mathrm{O_2^+}$ ($+\tfrac12$, met in Bartlett's $\mathrm{O_2^+[PtF_6]^-}$ below); in oxides, oxoanions and ordinary compounds it is negative (F is more electronegative than every other element including O). This restriction is a direct consequence of oxygen's very high electronegativity (second only to fluorine) and the first-row anomaly's coordination-number ceiling, which together prevent oxygen from being oxidized by anything except fluorine itself.

Catenation and allotropy differ sharply between O and S. Oxygen's stable allotrope is diatomic $\mathrm{O_2}$ (double bond, triplet ground state with two unpaired electrons — itself an unusual feature responsible for $\mathrm{O_2}$'s paramagnetism) with the strained, higher-energy $\mathrm{O_3}$ (ozone, bent, resonance-stabilized) as a reactive minor allotrope. Sulfur, by contrast, catenates extensively: the standard allotrope is $\mathrm{S_8}$, a puckered eight-membered ring held together by S–S single bonds, reflecting sulfur's (like phosphorus's) poor $3p$–$3p$ $\pi$-overlap and consequent preference for single bonds over the double-bonded $\mathrm{O_2}$ that efficient $2p$–$2p$ $\pi$-overlap gives oxygen. This is the same first-row-anomaly logic already used for $\mathrm{N_2}$ vs. $\mathrm{P_4}$ and $\mathrm{CO_2}$ vs. $\mathrm{SiO_2}$.

Key distinction. Do not describe sulfur's preference for $\mathrm{S_8}$ over a hypothetical $\mathrm{S_2}$ (analogous to $\mathrm{O_2}$) as "sulfur being less reactive." It is specifically about $\pi$-bonding efficiency: sulfur's larger $3p$ orbitals cannot form a strong $\pi$ bond, so single-bonded catenation wins out thermodynamically, exactly as for phosphorus.

2. Worked Example: Hydrogen Bonding and the H₂O/H₂S Boiling Point Anomaly

Problem. Explain why $\mathrm{H_2O}$ (boiling point $100\ \mathrm{^\circ C}$) boils at a far higher temperature than $\mathrm{H_2S}$ (boiling point $-60\ \mathrm{^\circ C}$), despite water having the lower molar mass (a factor that would, by itself, predict the opposite order for simple dispersion-force-dominated molecules).

Solution. Oxygen's very high electronegativity and small size concentrate a large partial negative charge on a compact atom, and each $\mathrm{O}$–$\mathrm{H}$ bond's hydrogen, bonded to this highly electronegative, lone-pair-rich oxygen, can form a strong, directional hydrogen bond to a neighbouring water molecule's lone pair. Sulfur, larger and less electronegative, supports no comparably strong S–H$\cdots$S hydrogen bonding network. The result is that liquid water is held together by an extensive three-dimensional hydrogen-bonded network requiring substantially more thermal energy to break apart (vaporize) than the comparatively weak dispersion forces between $\mathrm{H_2S}$ molecules — again, the first-row anomaly (oxygen's small size and electronegativity) is the root cause.

3. Group 17: The Unique Position of Fluorine

Group 17 (the halogens, $ns^2np^5$) shows the group oxidation state $+7$ (e.g. $\mathrm{ClO_4^-}$, perchlorate) for Cl, Br, and I, but fluorine shows only $-1$, never a positive oxidation state, because fluorine is the most electronegative element of all — there is simply no atom capable of oxidizing it. This is the sharpest, most absolute example of "first-row/lightest-member anomaly" behaviour in the entire p-block: it is not merely that fluorine prefers $-1$, as with oxygen's preference for $-2$, but that fluorine's oxidation state in every known compound is strictly $-1$ or $0$ (elemental $\mathrm{F_2}$).

The F–F bond is anomalously weak (bond enthalpy $\approx 158\ \mathrm{kJ\,mol^{-1}}$, noticeably weaker than Cl–Cl at $\approx 242\ \mathrm{kJ\,mol^{-1}}$, breaking the naive expectation that bond strength should simply fall monotonically down a group from the lightest, smallest-orbital member). The explanation is lone-pair repulsion: fluorine's very small atomic radius forces the three lone pairs on each F atom of $\mathrm{F_2}$ into very close proximity across the short F–F bond, and the resulting electron–electron repulsion destabilizes the bond significantly. This same small-size, lone-pair-repulsion effect also weakens the O–O single bond in $\mathrm{H_2O_2}$ relative to the S–S bond in $\mathrm{H_2S_2}$, and the N–N single bond in hydrazine relative to the P–P bond in diphosphane — a recurring, named exception to "bond strength falls down a group" that applies specifically to single bonds between two highly electronegative, lone-pair-rich, small period-2 atoms.

Oxidizing power falls down the group ($\mathrm{F_2} > \mathrm{Cl_2} > \mathrm{Br_2} > \mathrm{I_2}$), tracking decreasing electronegativity and, for $\mathrm{F_2}$ specifically, the weak F–F bond that makes fluorine's dissociation step unusually easy even though $\mathrm{F_2}$'s overall oxidizing strength (driven by the very favourable electron affinity and small ionic radius of $\mathrm{F^-}$, i.e. a large lattice/hydration-energy payoff) remains the largest in the group. This oxidizing-power trend directly underlies each halogen's ability to displace a lower halogen from its halide salt (e.g. $\mathrm{Cl_2} + 2\mathrm{Br^-} \rightarrow 2\mathrm{Cl^-} + \mathrm{Br_2}$, but not the reverse).

4. Hypervalency in the Halogens: Why Fluorine Cannot, but Iodine Can

Exactly as nitrogen cannot form $\mathrm{NCl_5}$ while phosphorus forms $\mathrm{PCl_5}$, fluorine cannot act as a central atom in a hypervalent interhalogen (there is no $\mathrm{FCl_3}$ or $\mathrm{FCl_5}$ with F central), while the larger halogens readily do: $\mathrm{ClF_3}$ (T-shaped, $Cl$ central), $\mathrm{ClF_5}$ (square pyramidal), $\mathrm{BrF_5}$, and the remarkable $\mathrm{IF_7}$ (pentagonal bipyramidal, coordination number 7 — the highest coordination number among the simple main-group molecular fluorides met in this chapter). The ordering principle is the same first-row-anomaly logic throughout: the central atom of a hypervalent interhalogen must be the larger, less electronegative halogen, since it needs both the physical size to accommodate several substituents and (on traditional accounts) accessible higher-energy orbitals or hyperconjugative stabilization to support the expanded electron count; the terminal, substituent halogens are typically the smaller, more electronegative ones (most commonly F).

Worked example. Rationalize the structures of $\mathrm{ClF_3}$ and $\mathrm{IF_7}$ using VSEPR. $\mathrm{ClF_3}$: chlorine is the central atom with 3 bonding pairs (to F) and 2 lone pairs, giving 5 electron domains in a trigonal bipyramidal electron geometry with both lone pairs equatorial (minimizing the more repulsive lone pair–lone pair and lone pair–bond pair interactions at $90^\circ$), leaving a T-shaped molecular geometry. $\mathrm{IF_7}$: iodine, the largest and most polarizable halogen capable of hypervalency at this level, has 7 bonding pairs and no lone pairs, giving the pentagonal bipyramidal geometry — realizable only because iodine's large size physically accommodates seven fluorine substituents around it, a coordination number fluorine itself, or even chlorine, cannot support.

5. Oxoacid Strength Trends in the Halogens

The chlorine oxoacids $\mathrm{HOCl}$, $\mathrm{HClO_2}$, $\mathrm{HClO_3}$, $\mathrm{HClO_4}$ show systematically increasing acid strength with increasing number of oxygen atoms (increasing oxidation state of Cl): $\mathrm{HOCl} \ll \mathrm{HClO_2} < \mathrm{HClO_3} < \mathrm{HClO_4}$ (the strongest common oxoacid known). Two reinforcing explanations are standard: (i) each additional, highly electronegative oxygen atom withdraws electron density inductively from the O–H bond, weakening it and increasing the ease of proton loss; (ii) the conjugate base's negative charge is increasingly delocalized (stabilized) over a greater number of oxygen atoms as the oxoanion's own oxidation state rises, lowering the base's affinity for the departed proton. This same two-part reasoning — inductive withdrawal plus increasing conjugate-base charge delocalization — explains the analogous trend among the oxoacids of any p-block nonmetal (nitrogen: $\mathrm{HNO_2} < \mathrm{HNO_3}$; sulfur: $\mathrm{H_2SO_3} < \mathrm{H_2SO_4}$) and will reappear quantitatively in the analytical chemistry block later in this course.

6. Group 18: The Fall of "Inert"

Until 1962, the noble gases (group 18, $ns^2np^6$, a filled valence shell) were assumed to be entirely unreactive — the very name "inert gases" encoded this assumption. Neil Bartlett overturned it by first preparing $\mathrm{O_2^+[PtF_6]^-}$ (noting that the ionization energy of $\mathrm{O_2}$, $1165\ \mathrm{kJ\,mol^{-1}}$, was closely matched by $\mathrm{PtF_6}$'s oxidizing power) and then reasoning that since $IE_1(\mathrm{Xe}) = 1170\ \mathrm{kJ\,mol^{-1}}$ is almost identical, xenon should be oxidizable by $\mathrm{PtF_6}$ too — and it was, giving the first noble-gas compound, $\mathrm{Xe[PtF_6]}$.

The key periodic insight is that noble-gas reactivity correlates directly with ionization energy, exactly as it does for every other element in this course: He, Ne, and Ar have prohibitively high $IE_1$ values and no known stable compounds (beyond a handful of extreme-matrix-isolation curiosities); Xe (and to a lesser extent Kr) has a comparatively modest $IE_1$, made further accessible by combination with the strongest oxidizers available (F, O). Xenon chemistry is now well developed: $\mathrm{XeF_2}$, $\mathrm{XeF_4}$, $\mathrm{XeF_6}$, $\mathrm{XeO_3}$, $\mathrm{XeOF_4}$, and the perxenate ion $\mathrm{XeO_6^{4-}}$ are all well-characterized, thermodynamically stable (if sometimes kinetically explosive/oxidizing) compounds.

Worked example — predicting $\mathrm{XeF_2}$'s shape. Xenon in $\mathrm{XeF_2}$ has 2 bonding pairs (to F) and 3 lone pairs, for 5 electron domains: trigonal bipyramidal electron geometry with all three lone pairs equatorial (again minimizing $90^\circ$ lone-pair interactions), giving a linear molecular geometry — consistent with the experimentally observed structure. Bonding in $\mathrm{XeF_2}$ is often described using a three-centre four-electron (3c-4e) bond along the F–Xe–F axis: unlike boron's electron-deficient 3c-2e bonds (too few electrons for the bonding framework), xenon's hypervalent bonding is electron-rich (more electrons than a simple two-centre picture can accommodate without invoking $d$-orbital participation), and the 3c-4e model distributes this excess density over a delocalized molecular orbital system spanning all three atoms, avoiding the need to expand xenon's octet in a literal sense.

Common Exam Pitfalls

• Claiming oxygen "never" has oxidation states other than $-2$. It is $-1$ in peroxides, $-\tfrac{1}{2}$ in superoxides, $0$ in $\mathrm{O_2}/\mathrm{O_3}$, and positive only in fluorides such as $\mathrm{OF_2}$ ($+2$) and $\mathrm{O_2F_2}$ ($+1$), or in $\mathrm{O_2^+}$ ($+\tfrac12$).
• Assuming bond strength falls monotonically down every group. The F–F, O–O (single), and N–N (single) bonds are all anomalously weak relative to their heavier congeners' analogous bonds, due to lone-pair repulsion across the unusually short bond of the small, lone-pair-rich period-2 atom — the opposite of the "simple size trend" students often assume.
• Confusing fluorine's $-1$-only restriction with a mere "preference," as for oxygen's $-2$ preference. Fluorine's restriction is absolute (no known compound shows F in any oxidation state other than $-1$ or $0$); oxygen's is only a strong, near-universal preference with only a few exceptions (oxygen fluorides, $\mathrm{O_2^+}$).
• Putting fluorine at the centre of a hypervalent interhalogen. $\mathrm{FCl_3}$ does not exist; the central atom of any stable interhalogen is always the larger, less electronegative halogen (Cl, Br, or I), exactly mirroring the N-vs-P hypervalency restriction from the previous chapter.
• Treating the noble gases as uniformly unreactive. Only He, Ne, and (largely) Ar remain essentially unreactive; Xe and Kr form a well-characterized, thermodynamically respectable chemistry once their comparatively modest ionization energies are matched against sufficiently strong oxidizers like F or O.
• Explaining oxoacid strength trends by "more oxygens means more acidic" without the mechanism. Always cite both the inductive electron withdrawal from additional electronegative oxygens and the delocalization/stabilization of the resulting conjugate base's negative charge — citing only one of the two is an incomplete answer.

Summary

• Group 16: oxygen is essentially restricted to non-positive oxidation states (positive only in fluorides such as $\mathrm{OF_2}$, $+2$, and $\mathrm{O_2F_2}$, $+1$, and in $\mathrm{O_2^+}$, $+\tfrac12$); $\mathrm{O_2}$ (double-bonded, paramagnetic) contrasts with sulfur's single-bonded, catenated $\mathrm{S_8}$ ring, mirroring the $\mathrm{N_2}$-vs-$\mathrm{P_4}$ first-row anomaly; water's exceptionally high boiling point (vs. $\mathrm{H_2S}$) arises from extensive O–H$\cdots$O hydrogen bonding.
• Group 17: fluorine shows only $-1$ (never positive); the F–F bond is anomalously weak from lone-pair repulsion across its short bond length; oxidizing power falls $\mathrm{F_2} > \mathrm{Cl_2} > \mathrm{Br_2} > \mathrm{I_2}$; hypervalent interhalogens ($\mathrm{ClF_3}$, $\mathrm{IF_7}$) always have the larger, less electronegative halogen as the central atom, never fluorine.
• Oxoacid strength increases with the number of oxygens/oxidation state of the central atom (e.g. $\mathrm{HOCl} < \mathrm{HClO_4}$), via combined inductive withdrawal and conjugate-base charge delocalization.
• Group 18: noble-gas reactivity tracks ionization energy exactly as elsewhere in the periodic table; Bartlett's 1962 preparation of $\mathrm{Xe[PtF_6]}$ overturned the "totally inert" assumption; xenon fluorides/oxides (e.g. $\mathrm{XeF_2}$, linear via VSEPR) are rationalized with hypervalent, electron-rich 3c-4e bonding, structurally and electronically the opposite extreme from boron's electron-deficient 3c-2e bonding.