A ladder of oxidation states, one rung apart
By now you have met nitrogen at its two extremes: as ammonia, NH3, where it sits at [[oxidation-state|oxidation state]] -3 having grabbed three electrons' worth of bonds from less electronegative hydrogen, and as the famously aloof N2 of the Haber-Bosch story, formally at 0. What makes nitrogen's chemistry so rich is everything *between* and *above* those two. With five valence electrons (2s2 2p3) and a small, electronegative core, nitrogen can comfortably exist at -3, -2, -1, 0, +1, +2, +3, +4, and +5 — almost a full house of integer states, and oxygen is the partner that walks it up the positive half of the ladder.
nitrogen's oxidation-state ladder (oxygen side) +5 N2O5 HNO3 / nitrate NO3- <- top of the ladder +4 NO2 , N2O4 +3 N2O3 HNO2 / nitrite NO2- +2 NO +1 N2O (laughing gas) 0 N2 <- the inert anchor -3 NH3 / ammonium NH4+ (the other end) oxide and oxoacid at the SAME rung are close cousins: N2O3 + H2O -> 2 HNO2 (+3) N2O5 + H2O -> 2 HNO3 (+5)
The nitrogen oxides: odd electrons and a brown haze
Start in the middle of the ladder, where the [[nitrogen-oxides|nitrogen oxides]] live. Nitric oxide, NO (+2), and nitrogen dioxide, NO2 (+4), share a strange trait you first met in the bonding rung: both have an *odd* total number of electrons, making them [[odd-electron-molecules|radicals]] with one unpaired electron. NO has 11 valence electrons; you simply cannot pair them all, so its molecular-orbital picture leaves one lonely electron in an antibonding pi orbital. This is why no neat octet Lewis structure ever quite satisfies you for these molecules — and why they are so reactive. Radicals are chemistry's restless singles, always looking to pair up.
Because NO2 is a radical, it does what restless singles do: it pairs up. Two NO2 molecules join their lonely electrons into one bond, giving the colourless dimer N2O4 (dinitrogen tetroxide). The two exist in a temperature-sensitive equilibrium, 2 NO2 (brown) <=> N2O4 (colourless). Warm a sealed tube and it darkens to chocolate brown as NO2 wins; chill it and the colour drains as N2O4 takes over — a vivid lecture-demo of Le Chatelier. That brown NO2 is exactly the colour you see hanging over a traffic-choked city: nitrogen oxides, collectively written NOx, are the breath of combustion engines, born when the furnace heat of an engine forces atmospheric N2 and O2 to react that would otherwise ignore each other entirely.
At the bottom of the positive ladder sits the gentle one: nitrous oxide, N2O (+1), better known as laughing gas. It is linear (N-N-O), surprisingly unreactive at room temperature, a mild anaesthetic, a whipped-cream propellant — and, less charmingly, a greenhouse gas roughly 300 times as potent per molecule as CO2, much of it released from over-fertilised farm soils. Notice the spread: the same two elements, nitrogen and oxygen, give a sleepy anaesthetic at +1, a colourless radical at +2, and a choking brown pollutant at +4. The oxidation state is not a mere label here — it genuinely sets the molecule's personality.
Smog, acid rain, and the atmosphere's two faces
Follow NOx out of the tailpipe and into the air, and you find nitrogen's redox restlessness doing real damage. In sunlight, NO2 photodissociates — light snaps it into NO plus a free oxygen atom — and that loose O atom seizes an O2 molecule to make ozone, O3. Down at street level ozone is a lung irritant, the sharp bite of photochemical smog. The same NOx, this is the cruel irony, also helps generate the haze, then helps generate the ozone that scours it, in a tangle of light-driven radical reactions over a sweltering city. The chemistry that warms an engine ends up rewriting the air a whole city breathes.
Higher up the ladder, NO2 dissolves and oxidises in cloud droplets all the way to nitric acid, HNO3, one of the two main culprits of acid rain (sulfuric acid, from sulfur dioxide, is the other — you will meet it in the chalcogen guides). Acidified rain leaches nutrients and aluminium from soils, acidifies lakes, and eats the calcium carbonate of statues and cathedrals. The remedy in cars is the [[catalytic-converter|catalytic converter]], a clever piece of inorganic catalysis: over a platinum-rhodium surface it reduces NOx back to harmless N2 while oxidising unburnt fuel and CO. It is, in effect, a small reactor bolted under the car whose whole job is to march nitrogen back down the oxidation ladder it was forced up in the engine.
Two oxoacids: nitrous (+3) and nitric (+5)
The two famous [[oxoacids-of-nitrogen|oxoacids of nitrogen]] sit on the +3 and +5 rungs. Nitrous acid, HNO2 (+3), is weak and unstable — you cannot bottle it; you make it cold, in solution, just before you need it, usually by acidifying a nitrite salt. Because nitrogen at +3 sits between the +5 of nitrate and the lower oxides, nitrous acid is a chemical fence-sitter: it can act as either an oxidant or a reductant depending on its partner, and it readily undergoes [[disproportionation|disproportionation]], where one element splits into a higher and a lower state at once: 3 HNO2 -> HNO3 + 2 NO + H2O. There, +3 nitrogen falls apart into +5 (in HNO3) and +2 (in NO) simultaneously — a single oxidation state shoving some of itself up and some of itself down.
Nitric acid, HNO3 (+5), is the heavyweight: a strong acid, fully dissociating in water to give H+ and the nitrate ion, NO3-. The nitrate ion is a small triumph of resonance — a flat, trigonal-planar ion in which the negative charge and the pi bonding are smeared evenly over all three equivalent N-O bonds, so the three bonds are identical and intermediate between single and double. That symmetric delocalisation is why nitrate is so stable and so abundant. But concentrated nitric acid wears a second hat far more dangerous than its acidity: it is a powerful [[oxidizing-and-reducing-agents|oxidising agent]]. It dissolves copper and silver — metals that mere strong acids like HCl cannot touch — not by the usual hydrogen-releasing route but by the nitrogen itself being reduced, pulling electrons out of the metal as it falls back down the ladder to NO2 or NO.
Two famous quirks follow from that oxidising power. Mix concentrated nitric and hydrochloric acids three-to-one and you get aqua regia, 'royal water', which dissolves even gold and platinum — the nitric acid oxidises the metal while the chloride ions clamp onto the resulting cation as chloride complexes, pulling the reaction forward. Yet, paradoxically, very concentrated nitric acid does *not* dissolve iron or aluminium: it oxidises their surface so fiercely and so fast that it grows a thin, dense, protective oxide skin — passivation — that seals the metal off. The same reagent that devours copper is quietly repelled by a sheet of aluminium. Reactivity, once again, is full of honest surprises.
The Ostwald process: ammonia in, nitric acid out
Where does industrial nitric acid come from? From ammonia — and so, ultimately, from the air. The [[ostwald-process|Ostwald process]] is the second half of a great two-stage human conquest of the nitrogen cycle: Haber-Bosch first pins inert N2 down into ammonia at -3, and Ostwald then walks that nitrogen all the way up the ladder to +5 in nitric acid. It is a tidy demonstration that you can climb the whole oxidation ladder on purpose if you have the right catalysts and conditions. Run the process backwards in your mind and it is just the oxidation of ammonia by oxygen, staged carefully so it stops at nitric acid rather than running away to plain N2.
- Burn ammonia in air over a hot platinum-rhodium gauze, around 850 C, for a fraction of a second: 4 NH3 + 5 O2 -> 4 NO + 6 H2O. The catalyst and the very short contact time are the trick — they coax nitrogen up to +2 (NO) and stop it short of N2.
- Cool the gases and let the NO meet more oxygen: 2 NO + O2 -> 2 NO2. Nitrogen has now climbed from +2 to +4. This step likes lower temperatures, so cooling actually helps it along.
- Absorb the NO2 in water with still more oxygen: 4 NO2 + O2 + 2 H2O -> 4 HNO3. Nitrogen reaches the top rung, +5, and out comes nitric acid — typically as a 60-some-percent solution ready for concentrating.
Watch the climb across the three steps: nitrogen goes -3 (ammonia) -> +2 (NO) -> +4 (NO2) -> +5 (HNO3). Every step is an oxidation, oxygen doing the lifting each time, and a clever cascade of temperatures and a precious-metal catalyst keeps the reaction marching up to the top rung instead of overshooting. It is one of the most elegant industrial sequences ever devised, and it is pure oxidation-state bookkeeping made flesh in stainless-steel towers.
Fertilizer and explosive: the nitrate's double life
Why pour all that industrial effort into making nitric acid? Because what you do with it feeds and arms the world. React nitric acid with ammonia and you get ammonium nitrate, NH4NO3 — a single salt that carries nitrogen in both of its useful biological forms (the +5 nitrate that plants take up, and the -3 ammonium). It is the workhorse fertilizer behind much of modern agriculture; a huge fraction of the protein in human bodies today contains nitrogen atoms that passed through an Ostwald tower. Strip nitrogen fertilizer out of the world and crop yields would collapse. This is the quiet, world-feeding face of nitrogen's redox chemistry.
But that very same molecule has a violent twin. Ammonium nitrate carries oxidising nitrate (+5) and reducible ammonium (-3) in one crystal — fuel and oxidiser packed into a single compound. Set it off and the two halves rush to meet in the middle, collapsing toward stable N2 and water and releasing a vast volume of hot gas almost instantly: that is what an explosive is. NH4NO3 is both a top fertilizer and a notorious blasting agent (and the cause of several catastrophic accidental detonations). The same logic underlies nitroglycerine and TNT: cram a fuel and a high-oxidation-state nitro group close together, and detonation is the thermodynamically enormous, kinetically sudden plunge back toward the inert N2 that nitrogen always secretly wants to be.
Step back and the whole guide tells one story. Nitrogen's enormous range of oxidation states is precisely what makes it so useful and so dangerous: it stores chemical energy by being held, against its will, at a high state far from the deep thermodynamic well of N2. The triple bond of N2 is one of the strongest in chemistry, so any time nitrogen can tumble back down to it, a great deal of energy is freed — gently in a catalytic converter, productively in a fertilizer factory, and catastrophically in an explosion. Master the ladder of oxidation states, and the smog, the rain, the laughing gas, the fertilizer, and the bomb all turn out to be the same chemistry, just read at different speeds.