Nitrogen and Phosphorus Compounds

01 / UNDERSTAND & APPLY

1. Start from the valence shell, then inspect the actual compound

Group 15 contains N, P, As, Sb and Bi among its familiar elements. Their outer-shell pattern is ns²np³. This common starting point explains recurring three-coordinate compounds and oxidation states such as −3, +3 and +5, but it does not make all their compounds structurally identical. Oxidation number is a formal electron-accounting label, not the number of bonds and not a claim that a free ion with that charge exists.

For example, nitrogen is +5 in nitrate but has only three oxygen neighbours; resonance descriptions keep nitrogen within its second-period octet. Ammonium has four N–H bonds and nitrogen remains −3. A statement such as “nitrogen cannot reach +5 because it cannot expand its octet” confuses oxidation state with coordination. Nitrate is the immediate counterexample.

Nitrogen is small and forms effective multiple bonds using its compact 2p orbitals. The very strong bond in N₂ makes activation difficult. Phosphorus more readily builds single-bonded frameworks and high-coordinate compounds. Describing PCl₅ as trigonal bipyramidal is a geometry statement; a literal promotion into five equivalent d-containing hybrids is not needed as a modern explanation of its bonding.

As atoms get larger down the group, metallic character generally increases. Nitrogen and phosphorus are nonmetals, arsenic and antimony show metalloid behaviour, and bismuth is metallic. Radius and ionization-energy trends are useful broad guides, but electron affinities and melting points are not simple monotonic ladders: electronic structure, allotropes and solid-state bonding matter.

02 / UNDERSTAND & APPLY

2. Dinitrogen, ammonia and the difference between rate and yield

N₂ has a short, strong triple bond. Its relative inertness under ordinary conditions is a kinetic observation: a substantial barrier must be overcome to break or strongly activate that bond. It does not mean nitrogen compounds are impossible or that every reaction involving nitrogen is thermodynamically unfavourable. Reactive metals can form nitrides; for example, 6Li+N₂→2Li₃N is a simple balanced representation.

Industrial ammonia synthesis uses N₂+3H₂⇌2NH₃, an exothermic equilibrium. Increasing pressure at fixed temperature favours the side with fewer gas molecules. Lower temperature favours equilibrium ammonia yield but slows the approach to equilibrium. A suitable iron-based catalyst provides a faster pathway without changing K at that temperature. Practical processing balances equilibrium, rate, separation and recycling rather than maximizing one variable in isolation.

An NH₃ molecule has three N–H bonds and one lone pair: approximately tetrahedral electron-domain arrangement, trigonal-pyramidal molecular shape and an H–N–H angle near 107°. Its polarity and hydrogen bonding help explain its strong interaction with water. Nitrogen’s lone pair can accept a proton, giving tetrahedral NH₄⁺. The four bonds in the resulting ammonium ion are equivalent; one does not remain permanently distinguishable as a special coordinate bond.

In water, NH₃+H₂O⇌NH₄⁺+OH⁻ is an equilibrium, so ammonia is a weak base rather than a source of a fixed one mole of free OH⁻ per mole dissolved. Reaction with a strong acid can nevertheless consume ammonia essentially stoichiometrically under suitable conditions. Its lone pair also binds metal ions, as in ammine complexes; acid–base protonation and coordination are related electron-pair-donation ideas, but their equilibria must be written separately.

03 / UNDERSTAND & APPLY

3. Nitrogen oxides: electron count, colour and acid anhydrides

The familiar sequence N₂O, NO, N₂O₃, NO₂ and N₂O₅ gives average nitrogen oxidation states +1, +2, +3, +4 and +5. An average is especially important for N₂O: its two nitrogen positions are inequivalent, so an average +1 does not prove identical atom-specific environments. N₂O is linear; NO is diatomic; NO₂ is bent. Formula patterns alone cannot supply those shapes.

NO has an odd electron and is paramagnetic, yet the ordinary gas is colourless. NO₂ also has an odd electron and is paramagnetic, but it is brown. Therefore paramagnetism does not imply visible colour. Two NO₂ molecules can form N₂O₄, pairing the odd electrons in a new bond; the dimer is diamagnetic and much less coloured. The gas equilibrium responds to temperature and pressure.

Dimer formation 2NO₂⇌N₂O₄ is exothermic. Cooling favours dimerization, while compression at fixed temperature favours the smaller gas-particle count. Heating increases the equilibrium fraction of NO₂. Colour intensity also depends on concentration and optical path length, so a colour observation needs a specified comparison rather than an unsupported claim about K.

N₂O and NO are classed as neutral oxides in the usual acid–base classification; “neutral” does not mean chemically inert or incapable of redox reactions. N₂O₃ is the formal anhydride of nitrous acid, while N₂O₅ is the anhydride of nitric acid. N₂O₃ is unstable toward NO/NO₂ at ordinary warmer conditions, so do not assume a bottle of pure stable room-temperature N₂O₃ gas. NO₂ reacts with water by disproportionation rather than producing an ordinary stable +4 nitrogen oxoacid.

04 / UNDERSTAND & APPLY

4. Nitric acid combines strong acidity with oxidizing chemistry

The conceptual Ostwald sequence first oxidizes ammonia catalytically to NO, then oxidizes NO to NO₂, then absorbs nitrogen oxides to form nitric acid with appropriate recycling and further oxidation. Keep each balanced stage separate: 4NH₃+5O₂→4NO+6H₂O; 2NO+O₂→2NO₂; 3NO₂+H₂O→2HNO₃+NO. These are reaction relationships, not instructions for operating a gas-handling process.

In ordinary dilute aqueous acid–base calculations, HNO₃ is a strong monoprotic acid. In metal reactions, nitrate can act as an oxidant: the gas need not be H₂. Copper gives NO as the principal textbook product with dilute nitric acid and NO₂ with concentrated nitric acid. Metal identity, concentration and other conditions affect nitrate reduction; no universal rule says that every dilute-acid metal reaction produces only NO.

The half-reaction NO₃⁻+4H⁺+3e⁻→NO+2H₂O requires three electrons per nitrate nitrogen. The NO₂ version NO₃⁻+2H⁺+e⁻→NO₂+H₂O requires one. Combining with Cu→Cu²⁺+2e⁻ explains different gas-to-metal ratios without memorizing coefficients independently. Extra nitrate remains as the counterion of the dissolved copper salt.

Concentrated nitric acid can passivate metals such as iron, aluminium and chromium through a protective surface film under appropriate conditions. Passivation is a kinetic surface effect, not proof that the metal has become intrinsically nonreactive. Conditions can disrupt the film. Avoid applying a simple activity-series rule without checking the oxidizing acid and the surface state.

05 / UNDERSTAND & APPLY

5. Nitrites and nitrates: medium and counterion change the outcome

Nitrite contains nitrogen at +3, so it can be oxidized to nitrate or reduced to lower oxidation states depending on its reaction partner. Nitrous acid is weak and unstable as an isolated ordinary aqueous solute; the balance 3HNO₂→HNO₃+2NO+H₂O illustrates disproportionation from +3 to +5 and +2. This explains why acidifying a nitrite cannot always be treated as simple permanent accumulation of HNO₂.

Nitrate contains nitrogen at +5, the familiar upper oxidation state. In ordinary redox questions it is reduced rather than further oxidized at nitrogen. The reduction product must still be specified or inferred from justified conditions. Merely assigning +5 is not enough to calculate an equivalent mass: nitrate-to-NO involves three electrons, whereas nitrate-to-NO₂ involves one.

Thermal decomposition of metal nitrates depends on the cation. The common alkali-metal pattern, except lithium, is nitrate to nitrite plus oxygen: 2NaNO₃→2NaNO₂+O₂. Lithium and many less electropositive metal nitrates give oxide, NO₂ and O₂; for example, 4LiNO₃→2Li₂O+4NO₂+O₂. Silver nitrate gives metal, NO₂ and oxygen in the elementary final-product model. These are theoretical product patterns, not safe student heating procedures.

The distinction between nitrate and nitrite is therefore chemical as well as spelling: oxygen count, oxidation state, acid strength and redox role differ. Never identify a gas by smelling it or infer safety from a colourless appearance. Toxic oxides and potentially hazardous nitrate decompositions belong to conceptual reaction analysis here.

06 / UNDERSTAND & APPLY

6. Phosphorus: molecule, network and oxide cages

White phosphorus consists of discrete P₄ tetrahedra. Each phosphorus atom bonds to three other phosphorus atoms and has a lone pair. The triangular faces impose 60° P–P–P angles, quite different from the relaxed geometry of an ordinary three-coordinate phosphorus centre. The strained small cage helps explain its high reactivity. A tetrahedral arrangement of four atom positions does not mean each P atom has four bonded neighbours.

Red phosphorus has extended covalent structures rather than independent P₄ molecules; black phosphorus has puckered layered bonding. Differences in connectivity explain substantial differences in physical properties and reactivity. These allotropes share elemental composition but not molecularity or crystal structure. White phosphorus is hazardous and highly reactive; no preparation or handling procedure is implied by the structural comparison.

The oxide formulas P₄O₆ and P₄O₁₀ are more informative molecular formulas than their empirical ratios P₂O₃ and P₂O₅. In the cage description, insert an oxygen into each of the six P–P edges of a P₄ tetrahedron to obtain six bridging oxygens in P₄O₆. Adding one terminal oxygen at each of four phosphorus sites gives P₄O₁₀. The oxygen counts follow from the topology.

Phosphorus(III) oxide forms phosphorous acid on hydrolysis; phosphorus(V) oxide forms phosphoric acid with sufficient water. P₄O₁₀ has a strong affinity for water and can act as a dehydrating agent. Its high phosphorus oxidation state does not automatically make it a powerful oxidant: stable P–O bonding matters. A formal P=O drawing is useful connectivity bookkeeping, but does not require a literal localized dπ–pπ bonding story.

07 / UNDERSTAND & APPLY

7. Hydrides and phosphorus chlorides require structure and phase labels

The familiar group hydrides EH₃ have three E–H bonds and a lone pair, giving pyramidal molecules. NH₃ is a substantially stronger base in water than PH₃. A lone pair’s presence alone does not set base strength: orbital character, bond formation and solvation all contribute. Heavier hydrides generally have weaker E–H bonds, lower thermal stability and stronger reducing tendency down the group. NH₃ also has unusually strong intermolecular hydrogen bonding compared with the heavier hydrides.

PCl₃ is pyramidal with a lone pair on phosphorus. Molecular gaseous PCl₅ is trigonal bipyramidal, with three equatorial and two axial positions. Its axial bonds are ordinarily longer and more readily affected than the equatorial bonds; an axial site has three 90° interactions with equatorial directions. This geometry does not apply unchanged to the ordinary ionic solid.

Solid PCl₅ is described by [PCl₄]⁺[PCl₆]⁻: tetrahedral cations and octahedral anions. The pair contains two P and ten Cl atoms, matching two PCl₅ formula units. Both phosphorus centres still have oxidation state +5 despite the opposite ion charges. PCl₅(g)⇌PCl₃(g)+Cl₂(g) is also a gas-phase dissociation equilibrium, so temperature and pressure can affect vapour composition.

Hydrolysis is a replacement of chlorine-containing connections by oxygen-containing ones, not a reduction of phosphorus. PCl₃+3H₂O→H₃PO₃+3HCl keeps P at +3. Limited hydrolysis of PCl₅ can be represented as PCl₅+H₂O→POCl₃+2HCl; further hydrolysis gives H₃PO₄. With sufficient water the total balance is PCl₅+4H₂O→H₃PO₄+5HCl. The HCl is part of the acidic product mixture and must not be omitted in a neutralization calculation.

08 / UNDERSTAND & APPLY

8. Count P–OH for basicity and inspect P–H for reducing behaviour

Hypophosphorous acid H₃PO₂ is conventionally drawn as H₂P(=O)OH. It has one P–OH group and two P–H bonds. Phosphorous acid H₃PO₃ is HP(=O)(OH)₂, with two P–OH groups and one P–H bond. Orthophosphoric acid H₃PO₄ is P(=O)(OH)₃, with three P–OH groups and no P–H. These connectivities, not the leading H₃ in every formula, determine ordinary aqueous acid basicity.

Thus H₃PO₂ is monoprotic, H₃PO₃ diprotic and H₃PO₄ triprotic. A hydrogen bonded directly to phosphorus is not counted as an ordinary ionizable acidic proton in these aqueous neutralization questions. Successive deprotonations also have different equilibrium constants; calling H₃PO₄ triprotic does not mean three protons dissociate completely at every pH or that every experimental indicator reaches a third equivalence point cleanly.

For these familiar acids, P–H-containing structures are reducing, whereas orthophosphoric acid is not an ordinary reducing acid. P oxidation states are +1, +3 and +5 respectively. Oxidizing H₃PO₂ fully to H₃PO₄ transfers four electrons per phosphorus; oxidizing H₃PO₃ to H₃PO₄ transfers two. Acid equivalents and redox equivalents therefore answer different questions for the same reagent.

Condensation joins phosphate units through P–O–P bridges while losing water. Two H₃PO₄ units can give H₄P₂O₇+H₂O; pyrophosphoric acid has four P–OH groups and is formally tetraprotic. Metaphosphoric acids are commonly cyclic or polymeric (HPO₃)ₙ structures, not necessarily an isolated monomer with the empirical formula. Bridging oxygen bears no acidic H; count terminal OH groups in the actual structure.

09 / UNDERSTAND & APPLY

9. Extend the reasoning to arsenic, antimony and bismuth

The heavier members keep the ns²np³ pattern but increasingly favour +3 relative to +5. The inert-pair description summarizes the decreasing participation of the ns² pair in the higher oxidation state. It is a useful chemical trend, not a claim that the pair is absolutely unavailable in every bond. Ligand identity and thermodynamic conditions still matter; strongly electronegative fluorine can stabilize high oxidation states.

Arsenic forms both +3 and +5 compounds. AsCl₃ hydrolysis can be summarized as AsCl₃+3H₂O⇌H₃AsO₃+3HCl. Arsenous acid is commonly represented as As(OH)₃, unlike phosphorous acid HP(=O)(OH)₂. Matching empirical formulas H₃EO₃ therefore do not prove matching E–H/O–H connectivity or identical acid strength. Arsenic(III) compounds can be oxidized to arsenic(V), depending on the oxidant and medium.

Antimony lies between the more nonmetallic arsenic chemistry and bismuth’s more metallic behaviour. Sb(III) and Sb(V) are both important. Sb₂O₃ is amphoteric: acid and sufficiently basic media can dissolve it through different product species. It is better to state the medium and relevant antimony species than force every aqueous reaction into a single unqualified molecular formula.

Bismuth strongly favours +3 in ordinary aqueous chemistry. Bi₂O₃ is predominantly basic and reacts with acid: Bi₂O₃+6H⁺→2Bi³⁺+3H₂O. BiCl₃ can hydrolyze to insoluble BiOCl; added acid can suppress this hydrolysis. Bismuth(V) compounds such as bismuthate are oxidizing because reduction to Bi(III) is favoured. This contrasts with stable phosphate, showing why the same formal +5 label is not enough to predict oxidizing strength.

The overall oxide trend is toward increasing basic character down the group, but classification must name the actual oxide and oxidation state. N₂O and NO are neutral exceptions; phosphorus oxides are acidic, arsenic(III) oxide has amphoteric behaviour with substantial acidic character, antimony(III) oxide is amphoteric, and bismuth(III) oxide is basic. Do not replace this chemically qualified comparison with “all Group 15 oxides are acidic”.

Moscovium belongs below bismuth in the modern group, but its short-lived synthetic nuclei do not supply the ordinary bulk-compound chemistry used in these questions. Extending a simple table into confidently asserted practical Mc hydride or oxide properties would go beyond the evidence and the scope of this lesson.

10 / UNDERSTAND & APPLY

10. Solved practice and a final audit

Treat these as original conceptual and numerical exercises. State the reaction endpoint before counting equivalents, the phase before assigning a geometry, and the connectivity before counting acidic hydrogens. None of the gas or phosphorus chemistry is a laboratory handling instruction.

• Do not equate oxidation state, formal charge and coordination number.
• Do not count P–H as P–OH in aqueous acid basicity.
• Do not extend the NH₃ hydrogen-bond or basicity pattern unchanged to PH₃.
• Do not infer an oxide’s acidity or oxidizing power from oxidation number alone.