The p-Block Elements (Class 12) — NEET Notes

The four groups at a glance

The Class 12 p-block runs across four columns of remarkable diversity. Group 15 opens with the gas that fills 78 percent of the air and closes with a heavy metal — nitrogen, phosphorus, arsenic, antimony, bismuth. Group 16 contains the oxygen we breathe and the sulphur we burn — oxygen, sulphur, selenium, tellurium, polonium. Group 17 — the halogens — are the most reactive non-metals on the table, going from fluorine (pale yellow gas, oxidises water) to iodine (violet solid that sublimes). Group 18 closes the table with the noble gases, six monoatomic, closed-shell species whose chemistry was thought impossible until 1962.

The chapter rewards a single organising habit: read every property as a vertical trend down the group, then ask which member breaks the trend. Nitrogen breaks Group 15, oxygen breaks Group 16, fluorine breaks Group 17. The reasons are always the same four — small size, high electronegativity, high ionisation enthalpy, no valence-shell d-orbitals. Hold this scaffold in mind and the entire chapter becomes a single argument about size and orbital availability.

Group 15 — the nitrogen family

The members are nitrogen, phosphorus, arsenic, antimony and bismuth, with the synthetic moscovium below. The valence configuration ns²np³ hosts a fully filled s-subshell and an exactly half-filled p-subshell — a configuration so symmetric that ionisation enthalpies for the group exceed those of Group 14 in the same period. Going down the group: nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids, bismuth is a true metal. Atomic radius grows from N (70 pm) to Bi (148 pm), with the biggest jump between N and P. The covalent-radius increase from As to Bi is small because completely filled d (and 4f) orbitals screen the nuclear charge poorly in the heavier elements.

Three oxidation states matter — −3, +3 and +5. The −3 state shrinks down the group as the element becomes more metallic — bismuth hardly shows it at all. The +5 state falls in stability down the group via the inert pair effect; BiF₅ is the only well-characterised Bi(V) compound. The +3 state, conversely, gains stability down the group. Nitrogen also shows +1, +2, +4 oxidation states (with oxygen) but cannot form a pentahalide because it lacks d-orbitals to expand its octet. Phosphorus, by contrast, cheerfully forms PCl₅ and even PF₆⁻.

Reactivity with hydrogen gives the hydrides EH₃. Stability falls in the order NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃, mirrored by falling bond dissociation enthalpy and rising reducing power; BiH₃ is the strongest reducer of the lot. Basicity follows the same downward order. NH₃ is the outlier on boiling point — 240 K versus 185 K for PH₃ — because nitrogen's small size and high electronegativity allow N−H to hydrogen-bond in the liquid state, an option closed to the heavier hydrides. With oxygen the elements form E₂O₃ and E₂O₅; the higher oxide is always more acidic, and acidic character of E₂O₃ drops down the group from clearly acidic (N, P) to amphoteric (As, Sb) to basic (Bi₂O₃).

Dinitrogen, ammonia, oxides of nitrogen and nitric acid

Dinitrogen is industrially obtained by fractional distillation of liquid air; N₂ (b.p. 77 K) distils first, leaving liquid O₂ (b.p. 90 K). In the laboratory it is prepared by warming an aqueous solution of NH₄Cl with NaNO₂: NH₄Cl(aq) + NaNO₂(aq) → N₂(g) + 2H₂O + NaCl. Very pure N₂ comes from thermal decomposition of sodium or barium azide. Its triple bond (bond enthalpy 941 kJ mol⁻¹) makes it virtually inert at room temperature; this very inertness is the source of the chemistry that follows — turning N₂ into anything reactive is hard.

Ammonia itself is a colourless pungent gas, m.p. 198 K, b.p. 240 K. The NH₃ molecule is trigonal pyramidal — sp³ nitrogen with three N−H bond pairs and one lone pair, HNH angle 107.8°. The lone pair makes ammonia a Lewis base; it precipitates hydroxides from many metal-ion solutions and forms deep-blue [Cu(NH₃)₄]²⁺ with Cu²⁺, the basis of a classic detection test. Industrially, NH₃ is consumed by the fertiliser industry (ammonium nitrate, urea, ammonium sulphate) and by the nitric-acid plants that feed it.

Oxides of nitrogen exist in every oxidation state from +1 to +5 — a richness no other element matches with oxygen alone. N₂O (laughing gas, +1) is a colourless neutral gas made by heating ammonium nitrate. NO (+2) is colourless and neutral, prepared from sodium nitrite and an Fe(II) salt in dilute sulphuric acid; it is also the species formed in the Haber-Ostwald sequence's first oxidation step. NO₂ (+4) is brown and acidic; it dimerises to colourless N₂O₄ on cooling because NO₂ has an odd valence-electron count. N₂O₅ (+5) is a colourless acidic solid prepared by dehydrating HNO₃ with P₄O₁₀. NEET 2018 (Q.75) asked which oxide of nitrogen is not a common pollutant — N₂O₅, because it requires laboratory dehydration to form and does not appear from natural oxidation chains.

Concentrated HNO₃ is a strong oxidiser that attacks almost every metal except gold and platinum. NEET 2016 (Q.4) tested its action on copper: Cu + 4HNO₃(conc.) → Cu(NO₃)₂ + 2NO₂ + 2H₂O. With dilute HNO₃, copper gives NO instead. Zinc gives NO₂ with concentrated acid and N₂O with dilute. Some metals — Cr, Al — are passivated by an unreactive oxide film and do not dissolve. The brown ring test relies on Fe²⁺ reducing NO₃⁻ to NO, which forms the brown [Fe(H₂O)₅(NO)]²⁺ complex at the interface with conc. H₂SO₄.

Phosphorus — allotropes, phosphine and oxoacids

Phosphorus exists in three principal allotropes — white, red and black — interconvertible by heat and pressure. White phosphorus is the reactive twin: a translucent waxy solid that glows in the dark (chemiluminescence), insoluble in water but soluble in CS₂, and so reactive that it catches fire in air to give P₄O₁₀. Its structure is a discrete P₄ tetrahedron with P−P−P bond angle of just 60° — the angular strain is the reason for its reactivity. Red phosphorus comes from heating white P at 573 K in an inert atmosphere; it is a polymer of P₄ tetrahedra linked into chains, iron-grey, non-poisonous, odourless, far less reactive. Black phosphorus, made from red P at 803 K under pressure, is the most stable allotrope of all.

Phosphine (PH₃) is prepared by hydrolysing calcium phosphide (Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃) or by warming white phosphorus with concentrated NaOH under an inert CO₂ atmosphere — the same reaction that scrubs white P. Pure PH₃ is non-inflammable, but trace P₂H₄ or P₄ vapour makes it spontaneously inflammable, which is the basis of its use in Holmes' marine signals. PH₃ is weakly basic — it gives PH₄⁺ salts only with strong acids (HBr, HI), unlike NH₃ which protonates easily. The reason is poor sp³ mixing on phosphorus: the H−P−H angle in PH₃ is just 93.6° (almost pure-p bonding), and the lone pair sits largely in the 3s, making it less available for donation.

Phosphorus halides. PCl₃ is a colourless oily liquid prepared by passing dry chlorine over heated white phosphorus, P₄ + 6Cl₂ → 4PCl₃. It is pyramidal (sp³) and hydrolyses in moisture to H₃PO₃ + 3HCl. PCl₅ is a yellowish-white powder prepared from P₄ + 10Cl₂, and is trigonal bipyramidal in the gas phase — three equivalent equatorial P−Cl bonds at 202 pm and two longer axial bonds at 240 pm, because axial bond pairs feel more repulsion. On heating PCl₅ dissociates to PCl₃ + Cl₂. With water it gives POCl₃ first and then H₃PO₄.

Oxoacids of phosphorus are a NEET favourite because they hide a basicity trick. In every oxoacid phosphorus is sp³ hybridised with a tetrahedral geometry; each acid has at least one P=O bond and one P−OH bond. Acids in lower oxidation states (< +5) additionally contain either a P−P bond (as in hypophosphoric, H₄P₂O₆) or a P−H bond (as in hypophosphorous, H₃PO₂). Crucially, only P−OH bonds ionise — P−H bonds are not acidic. So H₃PO₂ is monobasic (one P−OH, two P−H), H₃PO₃ is dibasic (two P−OH, one P−H), and H₃PO₄ is tribasic (three P−OH). NEET 2016 (Q.17) tested exactly this: phosphinic acid (H₃PO₂) is monoprotic while phosphonic acid (H₃PO₃) is diprotic. Acids with P−H bonds (H₃PO₂, H₃PO₃) are also good reducing agents — they reduce AgNO₃ to metallic silver.

Group 16 — the oxygen family (chalcogens)

The members are oxygen, sulphur, selenium, tellurium and polonium, plus the synthetic livermorium. The valence configuration is ns²np⁴. Oxygen and sulphur are non-metals, selenium and tellurium metalloids, polonium a radioactive metal. Down the group, atomic and ionic radii increase, ionisation enthalpy decreases, and electronegativity falls — oxygen is second only to fluorine in electronegativity (3.5 on Pauling's scale). Curiously, oxygen has a less negative electron-gain enthalpy than sulphur because its 2p orbitals are so compact that an incoming electron faces strong interelectronic repulsion.

Oxidation states are −2, +2, +4, +6, with +4 and +6 being the common positive states. The +6 state loses stability down the group (inert pair again — SF₆ is rock-stable, but PoO₃ is not isolated). Oxygen typically shows only −2; exceptions are OF₂ (+2) and O₂F₂ (+1), the only positive oxidation states possible because fluorine is more electronegative than oxygen. Oxygen exists as diatomic O₂ with a double bond — a consequence of its small size enabling pπ–pπ overlap — while sulphur catenates extensively into puckered S₈ rings.

Acidic character of the hydrides H₂E rises down the group: H₂O < H₂S < H₂Se < H₂Te. This is because the E−H bond becomes longer and weaker as the chalcogen gets bigger — easier to ionise H⁺ from a weaker bond. Boiling points, however, do not follow molar mass: H₂O boils at 373 K, anomalously high because of hydrogen bonding, while H₂S (213 K), H₂Se (232 K) and H₂Te (269 K) follow the expected upward trend with mass. NEET 2022 (Q.57) tested exactly this anomaly — both statements about the simple mass-ordering of boiling points were marked incorrect because of H₂O's hydrogen bonding.

Dioxygen, ozone and sulphur allotropes

Dioxygen is prepared in the laboratory by heating chlorates, nitrates or permanganates with MnO₂ as catalyst — 2KClO₃ → 2KCl + 3O₂ — or by thermal decomposition of higher oxides like Ag₂O or HgO. Industrially, it is fractionally distilled from liquid air. O₂ is paramagnetic despite its even electron count (a Class-11 MO-theory result worth remembering), reacts directly with most metals and non-metals, and the resulting binary compounds are called oxides. NIOS classifies oxides into acidic (non-metals or high-O.S. metals: SO₂, CO₂, Cl₂O₇, Mn₂O₇), basic (most metallic oxides: Na₂O, CaO), amphoteric (ZnO, Al₂O₃, SnO, PbO) and neutral (CO, NO, N₂O).

Ozone is the angular sp²-hybridised allotrope of oxygen, a resonance hybrid in which both O−O bond lengths are identical at 128 pm and the bond angle is 117°. It is prepared by passing dry O₂ through a silent electrical discharge in a Siemens ozoniser (rubber and cork must be avoided because ozone corrodes them). The conversion is endothermic, ΔH° = +142 kJ mol⁻¹ for 3O₂ → 2O₃; this is why ozone is unstable and decomposes spontaneously. As an oxidiser it liberates nascent oxygen, O₃ → O₂ + O: it oxidises PbS to PbSO₄, Fe²⁺ to Fe³⁺, iodide to iodine. The ozone-iodide reaction in borate-buffered KI is the basis of quantitative O₃ estimation. Ozone "tails mercury" — it converts the convex meniscus of clean mercury into a trail of droplets that stick to glass via Hg₂O, removed by dilute acid wash. CFCs and NO from supersonic aircraft are the threats to the stratospheric ozone layer.

Sulphur allotropes. The yellow rhombic (α) sulphur is the stable form at room temperature, density 2.06 g cm⁻³, melting at 385.8 K. Crystallised by slow evaporation of S in CS₂. The monoclinic (β) sulphur is stable above 369 K — formed by melting α-sulphur in a dish, cooling till a crust forms, piercing two holes and pouring out the remaining liquid; amber needles of β-sulphur are left behind. Both contain puckered crown-shaped S₈ rings; the difference is in crystal packing. At elevated temperatures (~ 1000 K), gaseous S₂ dominates, and like O₂ it is paramagnetic with two unpaired electrons in π* orbitals. Other forms include amorphous (precipitated from S₂O₃²⁻ + acid) and plastic sulphur (pouring molten S into cold water; long chains that revert to S₈ rings on standing).

Sulphur dioxide SO₂ is produced when sulphur burns in air or in the lab by treating a sulphite with dilute H₂SO₄. Industrially it appears as a roasting by-product: 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂. SO₂ is a colourless pungent gas, highly soluble in water giving sulphurous acid H₂SO₃. The molecule is angular, a resonance hybrid in which both S−O bonds are equal. With water and chlorine in charcoal catalyst, it gives SO₂Cl₂ (sulphuryl chloride). With V₂O₅ and O₂ it goes to SO₃ — the cornerstone reaction of the Contact process. Moist SO₂ reduces Fe³⁺ to Fe²⁺ and decolourises acidified KMnO₄; the latter is the standard test for the gas. NEET 2022 (Q.100) asked which factors enhance SO₂ pollution — particulate matter (catalyses SO₂ → SO₃), ozone and H₂O₂.

Sulphuric acid — the king of chemicals

A nation's industrial strength is often measured in tonnes of sulphuric acid. Manufactured almost exclusively today by the Contact process, the route runs through three steps: produce SO₂, oxidise SO₂ to SO₃ with V₂O₅ catalyst, absorb SO₃ in concentrated H₂SO₄ to give oleum (H₂S₂O₇). Oleum is then carefully diluted with water to whatever concentration the customer needs. The plant yields 96–98 % pure H₂SO₄.

Sulphuric acid's chemistry comes from four characters: low volatility, strong acidity, dehydrating power, oxidising power. Ionisation is two-step — K_a1 > 10 (essentially complete first dissociation) but K_a2 = 1.2 × 10⁻². The acid liberates more volatile acids from their salts (2MX + H₂SO₄ → 2HX + M₂SO₄; X = F, Cl, NO₃). Its dehydrating action is brutal — it removes water from sugar to leave a black mass of carbon (C₁₂H₂₂O₁₁ → 12C + 11H₂O), and removes water of crystallisation from blue CuSO₄·5H₂O to leave white anhydrous CuSO₄. Hot concentrated H₂SO₄ is a moderately strong oxidiser — Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O; C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O.

Oxoacids of sulphur form a wide family that NEET tests almost annually. Sulphurous acid H₂SO₃ (two S−OH, one S=O), sulphuric acid H₂SO₄ (two S−OH, two S=O), thiosulphuric acid H₂S₂O₃ (one S=S bond), pyrosulphuric H₂S₂O₇ (one S−O−S bridge), peroxydisulphuric H₂S₂O₈ (one S−O−O−S peroxy bridge — Caro's twin), and the polythionates H₂S_nO₆ (chain of central S atoms with terminal SO₃⁻ groups). NEET 2023 (Q.96) matched these acids to their bond inventories — the key recognition pattern is whether the central linkage is S−O−S (pyro) or S−O−O−S (peroxy). NEET 2020 (Q.156) asked which oxoacid has a −O−O− linkage — answer: H₂S₂O₈. NEET 2017 (Q.13) asked which pair of ions contains an S−S bond: S₄O₆²⁻ and S₂O₃²⁻.

Group 17 — the halogen family

Members are fluorine, chlorine, bromine, iodine, astatine, plus synthetic tennessine. The valence configuration ns²np⁵ places each halogen one electron short of the next noble gas. They have the smallest atomic radii in their respective periods (maximum effective nuclear charge), the highest ionisation enthalpies among non-metals after the noble gases, the most negative electron-gain enthalpies of the periodic table, and the highest electronegativities. Fluorine is, on Pauling's scale, the most electronegative element of all. Halogens display a smooth physical gradation — F₂ and Cl₂ are gases (yellow, greenish-yellow), Br₂ is a red liquid, I₂ a violet solid; melting and boiling points rise steadily with atomic number. All halogens are coloured because they absorb in the visible range.

Anomalies of fluorine: ionisation enthalpy, electronegativity and electrode potential are all higher than the trends predict; ionic and covalent radii, m.p., b.p., bond dissociation enthalpy and electron-gain enthalpy are lower than expected. The cause is its uniquely small size, highest electronegativity, low F−F bond dissociation enthalpy (158 kJ mol⁻¹, lower than Cl₂!) and absence of valence-shell d-orbitals. The low F−F bond enthalpy comes from strong lone-pair–lone-pair repulsion in the tiny F₂ molecule. Fluorine shows only −1; all other halogens show +1, +3, +5, +7 in combination with O and F because they have d-orbitals to expand their octet.

Chlorine preparation. In the lab, MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O, or 2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂. Industrially, the Deacon process oxidises HCl with atmospheric O₂ over a CuCl₂ catalyst at 723 K: 4HCl + O₂ → 2Cl₂ + 2H₂O. Most chlorine, however, is obtained as a by-product of the electrolysis of brine in the chlor-alkali industry — Cl₂ at the anode, H₂ at the cathode, NaOH in solution. Cl₂ is a greenish-yellow gas, 2.5 times heavier than air. It reacts with metals (2Na + Cl₂ → 2NaCl), with hydrogen (H₂ + Cl₂ → 2HCl), and with cold dilute NaOH to give a chloride-hypochlorite mixture (the basis of bleaching powder: 2Ca(OH)₂ + 2Cl₂ → Ca(OCl)₂ + CaCl₂ + 2H₂O). Hot concentrated NaOH gives chloride + chlorate — a classic disproportionation in which Cl(0) becomes Cl(−1) and Cl(+5) simultaneously.

Hydrogen halides and oxoacids of halogens

HCl is prepared in the laboratory by heating sodium chloride with concentrated sulphuric acid at 420 K (NaCl + H₂SO₄ → NaHSO₄ + HCl); at 823 K the bisulphate reacts further to give Na₂SO₄. HCl gas is dried by passing through conc. H₂SO₄. It is colourless, pungent, freezes at 159 K, boils at 189 K, and is extremely soluble in water — its aqueous solution is hydrochloric acid, a strong acid by virtue of complete dissociation. With ammonia it gives white NH₄Cl fumes. Mixed 3 : 1 with concentrated HNO₃ it gives aqua regia, which dissolves gold and platinum.

Bond dissociation enthalpy falls H−F > H−Cl > H−Br > H−I, so stability falls in the same order. But acid strength rises HF < HCl < HBr < HI — because longer, weaker bonds release H⁺ more readily. HF is anomalously a liquid at room temperature (b.p. 293 K) because of strong F···H hydrogen bonding; HCl, HBr, HI are gases. HF also etches glass, attacking the SiO₂ network to give SiF₄ and water.

Oxoacids of chlorine form a tidy +1, +3, +5, +7 ladder — hypochlorous (HOCl), chlorous (HOClO), chloric (HOClO₂), perchloric (HOClO₃). Acid strength rises smoothly with oxidation state: HOCl < HOClO < HOClO₂ < HOClO₃. The mechanism is electron-pulling — more O atoms bonded to Cl means more electrons pulled from the O−H bond, weakening it. Perchloric acid is one of the strongest mineral acids known, requiring distillation under reduced pressure to handle. NEET 2016 (Q.32) tested exactly this order. The full halogen family: HOF is the only F oxoacid (fluorine cannot expand octet so cannot host +3 or higher); HOCl, HOClO, HOClO₂, HOClO₃ for Cl; HOBr, HOBrO₂, HOBrO₃ for Br; HOI, HOIO₂, HOIO₃ for I. Fluorine sits in HOF in the +1 oxidation state — the only positive oxidation state it ever takes, because oxygen is less electronegative than fluorine. NEET 2018 (Q.62) marked as bonus because the trap option "all halogens but fluorine show positive oxidation states" omits the HOF exception.

Interhalogen compounds

When two halogens combine they form interhalogen compounds of types XX', XX'₃, XX'₅, XX'₇, where X is the larger, less electronegative halogen and X' is the smaller, more electronegative one. The ratio of X to X' radii grows with the X' subscript, which is why only IF₇ exists in the X'₇ family — iodine is big enough and fluorine small enough to pack seven Fs around one I. Preparation is by direct combination (Cl₂ + F₂ → 2ClF; Cl₂ + 3F₂ → 2ClF₃; I₂ + 5F₂ → 2IF₅; I₂ + 7F₂ → 2IF₇) or by the action of a halogen on a lower interhalogen.

Shapes follow VSEPR cleanly: XX' is linear (sp³, three lone pairs); XX'₃ is bent T-shape (sp³d, two lone pairs in equatorial positions — BrF₃, ClF₃); XX'₅ is square pyramidal (sp³d², one lone pair — IF₅, BrF₅, ClF₅); XX'₇ is pentagonal bipyramidal (sp³d³, no lone pair — IF₇). NEET 2017 (Q.16) tested precisely this correspondence. Interhalogens are more reactive than the parent halogens except F₂, because the heteronuclear X−X' bond is weaker than the homonuclear X−X bond — NEET 2022 (Q.81) examined this with the assertion that ICl is more reactive than I₂ because the I−Cl bond is weaker than the I−I bond. All interhalogens hydrolyse to give the hydrohalic acid of the smaller halogen and an oxoacid of the larger one: ClF₅ + 3H₂O → 5HF + HClO₃.

Group 18 — the noble gases

The members are helium, neon, argon, krypton, xenon, radon, plus synthetic oganesson. All are monoatomic gases. Their valence configurations end in ns²np⁶ — a closed octet — except for He, which has 1s². This closed-shell structure is the source of their extreme inertness. Ionisation enthalpy is the highest in each period (He = 2372 kJ mol⁻¹; Xe = 1170 kJ mol⁻¹) and electron-gain enthalpy is large and positive (they have no appetite for an extra electron). Down the group the ionisation enthalpy falls smoothly as the atom gets bigger and the outer electrons less tightly bound. Atmospheric abundance: Ar is by far the most abundant (0.934 % of dry air), He is harvested from natural gas (some natural-gas wells contain up to 10 %), Xe and Kr are present in parts per million.

Xenon compounds. The first noble gas compound was prepared in 1962 by Neil Bartlett. He noticed that O₂'s first ionisation enthalpy (1175 kJ mol⁻¹) was almost identical to xenon's (1170 kJ mol⁻¹); since he had already made O₂⁺PtF₆⁻, he reasoned the same trick would work for Xe — and produced the red solid XePtF₆. Within months, direct fluorination gave the three binary fluorides:

• XeF₂ — Xe + F₂ at 673 K, 1 bar, Xe in excess. Linear molecule, sp³d hybridised with three lone pairs in the equatorial plane.
• XeF₄ — Xe + 2F₂ at 873 K, 7 bar, 1 : 5 ratio. Square planar, sp³d², two lone pairs trans.
• XeF₆ — Xe + 3F₂ at 573 K, 60–70 bar, 1 : 20 ratio. Distorted octahedral (sp³d³, one lone pair).

All three are colourless crystalline solids that sublime at room temperature, and all are powerful fluorinating agents. Hydrolysis distinguishes them clearly. XeF₂ hydrolyses slowly: 2XeF₂ + 2H₂O → 2Xe + 4HF + O₂. XeF₄ reacts violently to give a disproportionation: 6XeF₄ + 12H₂O → 4Xe + 2XeO₃ + 24HF + 3O₂. XeF₆ gives complete hydrolysis to XeO₃ in excess water, but partial hydrolysis gives oxofluorides: XeF₆ + H₂O → XeOF₄ + 2HF, and XeF₆ + 2H₂O → XeO₂F₂ + 4HF. XeO₃ is a colourless explosive solid with pyramidal geometry (sp³, one lone pair, three Xe=O bonds). XeO₄ is made by treating Ba₂XeO₆ with conc. H₂SO₄.

NEET 2016 (Q.12) tested the entire shape matrix in a single matching question: XeF₆ (distorted octahedral), XeO₃ (pyramidal), XeOF₄ (square pyramidal), XeF₄ (square planar). The xenon fluorides also act as fluoride donors or acceptors with Lewis acids and bases — XeF₂ + PF₅ → [XeF]⁺[PF₆]⁻; XeF₄ + SbF₅ → [XeF₃]⁺[SbF₆]⁻; XeF₆ + MF → M⁺[XeF₇]⁻.

Uses of noble gases: He fills weather balloons (non-flammable), cools superconducting magnets in MRI scanners (b.p. 4.2 K is the lowest of any substance), and dilutes O₂ in deep-sea diving cylinders (low solubility in blood prevents nitrogen narcosis). Ne fills discharge tubes for advertising signs (the red-orange glow), greenhouse and botanical lights. Ar provides the inert atmosphere in metallurgical arc welding and fills incandescent bulbs (prevents tungsten oxidation). Kr and Xe appear in specialised high-intensity lamps. Rn is radioactive and used (in trace amounts) in cancer radiotherapy.

The diagonal relationship is the subtle pattern that runs across both Class 11 and Class 12 p-block chemistry. The first member of each group resembles the second member of the next group diagonally below it — Li resembles Mg, Be resembles Al, B resembles Si — because as one moves rightward across a period electronegativity rises and atomic size falls, while moving down a group both rise. The two effects can cancel diagonally, giving similar charge density and similar chemistry. Within Class 12 p-block, the trace of this is N's pπ–pπ chemistry resembling C's more than P's, and O behaving like N at times in compactness-of-bonding.