Dinitrogen: Preparation & Properties
Ammonia and nitric acid both begin with dinitrogen, so a brief account of $\ce{N2}$ frames the whole topic. Commercially, dinitrogen is obtained by the liquefaction and fractional distillation of air; liquid dinitrogen (b.p. 77.2 K) distils off first, leaving behind liquid oxygen (b.p. 90 K). In the laboratory it is prepared by warming an aqueous solution of ammonium chloride with sodium nitrite.
$$\ce{NH4Cl(aq) + NaNO2(aq) -> N2(g) + 2H2O(l) + NaCl(aq)}$$
Small amounts of $\ce{NO}$ and $\ce{HNO3}$ accompany this reaction; they are removed by passing the gas through aqueous sulphuric acid containing potassium dichromate. Very pure nitrogen comes from the thermal decomposition of sodium or barium azide.
$$\ce{Ba(N3)2 ->[\Delta] Ba + 3N2} \qquad \ce{2NaN3 ->[\Delta] 2Na + 3N2}$$
Dinitrogen is a colourless, odourless, tasteless and non-toxic gas with very low water solubility. It is rather inert at room temperature because of the very high $\ce{N#N}$ bond enthalpy; reactivity rises sharply with temperature. At high temperatures it forms ionic nitrides with active metals and covalent nitrides with non-metals, and combines with hydrogen over a catalyst (Haber's process) and with dioxygen only near 2000 K.
| Reaction | Equation | Condition |
|---|---|---|
| With lithium | $\ce{6Li + N2 ->[\Delta] 2Li3N}$ | Heat; ionic nitride |
| With magnesium | $\ce{3Mg + N2 ->[\Delta] Mg3N2}$ | Heat; ionic nitride |
| With hydrogen | $\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$ | 773 K, catalyst (Haber) |
| With dioxygen | $\ce{N2 + O2(g) <=> 2NO(g)}$ | ~2000 K |
Ammonia: Preparation & Haber Process
Ammonia occurs in small quantities in air and soil, formed by the decay of nitrogenous organic matter such as urea. On a small scale it is obtained from ammonium salts by warming them with a strong base — caustic soda or calcium hydroxide.
$$\ce{2NH4Cl + Ca(OH)2 -> 2NH3 + 2H2O + CaCl2}$$ $$\ce{(NH4)2SO4 + 2NaOH -> 2NH3 + 2H2O + Na2SO4}$$
On the industrial scale, ammonia is manufactured by Haber's process, the direct combination of dinitrogen and dihydrogen. The reaction is reversible and exothermic.
$$\ce{N2(g) + 3H2(g) <=> 2NH3(g)} \qquad \Delta_f H^\circ = -46.1\ \text{kJ mol}^{-1}$$
The flow of the plant is shown below: a 1 : 3 mixture of purified $\ce{N2}$ and $\ce{H2}$ is compressed, passed over the heated catalyst, and the ammonia is condensed out while the unreacted gases are recycled.
Purified $\ce{N2}$ and $\ce{H2}$ in 1 : 3 ratio are compressed to about 200 atm, passed over a hot iron-oxide catalyst near 700 K, and the $\ce{NH3}$ is liquefied out while unconverted gases return to the loop.
Conditions for Maximum Yield
Because the forward reaction is exothermic and proceeds with a decrease in the number of gaseous molecules (4 mol of reactant gas → 2 mol of product), Le Chatelier's principle dictates that high pressure and low temperature favour ammonia. In practice the temperature cannot be too low or the rate becomes uselessly slow, so a compromise is struck.
| Factor | Optimum condition | Why |
|---|---|---|
| Pressure | ~200 atm ($200\times10^5$ Pa) | Shifts equilibrium toward fewer gas moles, i.e. toward $\ce{NH3}$ |
| Temperature | ~700 K | Compromise: low T favours yield, but moderate T is needed for a workable rate |
| Catalyst | Iron oxide | Speeds attainment of equilibrium (no shift in position) |
| Promoters | Small amounts of $\ce{K2O}$, $\ce{Al2O3}$ | Enhance catalyst efficiency |
The supplement notes that an earlier version of the process used iron as catalyst with molybdenum as the promoter; the modern catalyst is iron oxide carrying small quantities of $\ce{K2O}$ and $\ce{Al2O3}$.
A catalyst does not increase the yield
Students often write that the iron catalyst "increases the yield of ammonia." It does not. A catalyst only speeds the attainment of equilibrium; it leaves the equilibrium position — and hence the maximum yield — unchanged. Yield is governed by pressure and temperature.
Catalyst → faster equilibrium; pressure (high) and temperature (moderate, exothermic) → higher yield.
Structure of Ammonia
The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex. Nitrogen is $sp^3$ hybridised and carries three bond pairs and one lone pair; the lone pair pushes the three $\ce{N-H}$ bonds together, lowering the angle below the tetrahedral value. In the solid and liquid states ammonia molecules are associated through hydrogen bonds, just like water, which explains its melting and boiling points being higher than expected from its molecular mass.
Three bond pairs and one lone pair around $sp^3$ nitrogen give a pyramid; the lone pair is the source of ammonia's basic and ligand behaviour.
Properties & Reactions of Ammonia
Ammonia is a colourless gas with a pungent odour (f.p. 198.4 K, b.p. 239.7 K) and is highly soluble in water. Its chemistry is dominated by the lone pair on nitrogen, which makes it both a Brønsted base and a Lewis base.
As a base in water and with acids
In water the aqueous solution is weakly basic owing to $\ce{OH-}$ formation, and ammonia neutralises acids to give ammonium salts.
$$\ce{NH3(g) + H2O(l) <=> NH4+(aq) + OH^-(aq)}$$
Precipitating metal hydroxides
As a weak base it precipitates the hydroxides (or hydrated oxides) of many metals from their salt solutions — a behaviour exploited in qualitative analysis.
$$\ce{ZnSO4(aq) + 2NH4OH(aq) -> Zn(OH)2(s)\ (white) + (NH4)2SO4(aq)}$$ $$\ce{FeCl3(aq) + 3NH4OH(aq) -> Fe2O3.xH2O(s)\ (brown) + 3NH4Cl(aq)}$$
Ammonia's pyramidal shape and lone-pair chemistry sit inside the wider family trends — see Group 15: The Nitrogen Family.
Lewis base and complex formation
The lone pair lets ammonia bind metal ions, forming coordination complexes used to detect $\ce{Cu^2+}$ and $\ce{Ag+}$. With copper(II) ions the pale blue solution turns deep blue.
$$\ce{Cu^2+(aq)\ (blue) + 4NH3(aq) <=> [Cu(NH3)4]^2+(aq)\ (deep\ blue)}$$
The same donor ability dissolves silver chloride: the white precipitate $\ce{AgCl}$ redissolves in excess ammonia by forming the diamminesilver(I) complex.
$$\ce{AgCl(s)\ (white) + 2NH3(aq) -> [Ag(NH3)2]Cl(aq)}$$
Nitric Acid: Ostwald Process
Nitric acid is the most important of nitrogen's oxoacids ($\ce{H2N2O2}$ hyponitrous, $\ce{HNO2}$ nitrous, $\ce{HNO3}$ nitric). In the laboratory it is prepared by heating $\ce{KNO3}$ or $\ce{NaNO3}$ with concentrated sulphuric acid in a glass retort.
$$\ce{NaNO3 + H2SO4 -> NaHSO4 + HNO3}$$
On the large scale it is made by Ostwald's process, the catalytic oxidation of the ammonia produced by Haber's process. The sequence is three steps: oxidise $\ce{NH3}$ to $\ce{NO}$, oxidise $\ce{NO}$ to $\ce{NO2}$, then absorb $\ce{NO2}$ in water.
$$\ce{4NH3(g) + 5O2(g) ->[\text{Pt/Rh}][\text{500 K, 9 bar}] 4NO(g) + 6H2O(g)}$$ $$\ce{2NO(g) + O2(g) <=> 2NO2(g)}$$ $$\ce{3NO2(g) + H2O(l) -> 2HNO3(aq) + NO(g)}$$
The $\ce{NO}$ released in the last step is recycled. The aqueous acid is concentrated by distillation to about 68% by mass; further concentration to roughly 98% is achieved by dehydration with concentrated $\ce{H2SO4}$.
Left: the three-step Ostwald sequence with $\ce{NO}$ recycled. Right: in the gas phase $\ce{HNO3}$ is a planar molecule, with the two N–O bonds to terminal oxygens equivalent by resonance and one N–O(H) bond.
Structure & Acidic Nature
Pure nitric acid is a colourless liquid (f.p. 231.4 K, b.p. 355.6 K); laboratory grade contains about 68% $\ce{HNO3}$ by mass with a specific gravity of 1.504. In the gaseous state it exists as a planar molecule. In aqueous solution it behaves as a strong acid, ionising essentially completely to give hydronium and nitrate ions.
$$\ce{HNO3(aq) + H2O(l) -> H3O+(aq) + NO3^-(aq)}$$
HNO3 as an Oxidising Agent
The defining feature of nitric acid for NEET is that concentrated $\ce{HNO3}$ is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. Crucially, it does not generally liberate hydrogen; instead the nitrogen is reduced, and the product depends on the concentration of the acid, the temperature, and the nature of the metal. Copper is the classic example.
| Metal | Dilute HNO₃ | Concentrated HNO₃ |
|---|---|---|
| Copper | $\ce{3Cu + 8HNO3 -> 3Cu(NO3)2 + 2NO + 4H2O}$ gives NO |
$\ce{Cu + 4HNO3 -> Cu(NO3)2 + 2NO2 + 2H2O}$ gives brown NO₂ |
| Zinc | $\ce{4Zn + 10HNO3 -> 4Zn(NO3)2 + 5H2O + N2O}$ gives N₂O |
$\ce{Zn + 4HNO3 -> Zn(NO3)2 + 2H2O + 2NO2}$ gives NO₂ |
Concentrated nitric acid also oxidises several non-metals and their compounds: iodine to iodic acid, carbon to carbon dioxide, sulphur to sulphuric acid and phosphorus to phosphoric acid, each releasing $\ce{NO2}$.
$$\ce{I2 + 10HNO3 -> 2HIO3 + 10NO2 + 4H2O}$$ $$\ce{C + 4HNO3 -> CO2 + 2H2O + 4NO2}$$ $$\ce{S8 + 48HNO3 -> 8H2SO4 + 48NO2 + 16H2O}$$ $$\ce{P4 + 20HNO3 -> 4H3PO4 + 20NO2 + 4H2O}$$
Dilute vs concentrated — and the noble-metal exception
Do not assume the gas is always $\ce{NO2}$. With dilute $\ce{HNO3}$ copper gives colourless $\ce{NO}$; with concentrated acid it gives reddish-brown $\ce{NO2}$. Zinc gives $\ce{N2O}$ with dilute and $\ce{NO2}$ with concentrated. And note that gold and platinum resist nitric acid alone — they dissolve only in aqua regia.
Cu + dilute → NO · Cu + conc. → NO₂ · Zn + dilute → N₂O · Zn + conc. → NO₂.
Brown Ring Test & Passivity
Brown ring test for nitrate
The brown ring test identifies the nitrate ion. It relies on $\ce{Fe^2+}$ reducing nitrate to nitric oxide, which then binds another $\ce{Fe^2+}$ to give a brown complex. The test is performed by adding freshly prepared dilute ferrous sulphate to a solution containing nitrate, then pouring concentrated sulphuric acid carefully down the side of the tube; a brown ring forms at the interface of the two layers.
$$\ce{NO3^- + 3Fe^2+ + 4H+ -> NO + 3Fe^3+ + 2H2O}$$ $$\ce{[Fe(H2O)6]^2+ + NO -> [Fe(H2O)5(NO)]^2+\ (brown) + H2O}$$
Passivity of iron, chromium and aluminium
Some metals such as $\ce{Cr}$ and $\ce{Al}$ do not dissolve in concentrated nitric acid, because the acid oxidises the metal surface to a thin, adherent layer of oxide. This passive film seals the metal from further attack — the reason iron, chromium and aluminium become "passive" and appear unreactive toward concentrated $\ce{HNO3}$ despite its strength as an oxidiser.
Uses of Ammonia & Nitric Acid
| Compound | Major uses |
|---|---|
| Ammonia | Nitrogenous fertilisers (ammonium nitrate, urea, ammonium phosphate, ammonium sulphate); manufacture of inorganic nitrogen compounds, chiefly nitric acid; liquid $\ce{NH3}$ as a refrigerant. |
| Nitric acid | Manufacture of ammonium nitrate for fertilisers and nitrates for explosives and pyrotechnics; preparation of nitroglycerin, TNT and other organic nitro compounds; pickling of stainless steel; etching of metals; oxidiser in rocket fuels. |
| Dinitrogen | Manufacture of ammonia and other nitrogen chemicals (e.g. calcium cyanamide); inert atmosphere in the iron and steel industry; liquid $\ce{N2}$ as a refrigerant and in cryosurgery. |
Ammonia & nitric acid in one screen
- Haber: $\ce{N2 + 3H2 <=> 2NH3}$, exothermic; max yield at ~200 atm and ~700 K with Fe-oxide catalyst ($\ce{K2O}$, $\ce{Al2O3}$ promoters). Catalyst speeds equilibrium, not yield.
- NH₃ structure: trigonal pyramidal, $sp^3$, three bond pairs + one lone pair; the lone pair makes it a Lewis base.
- NH₃ reactions: weakly basic in water; precipitates $\ce{Zn(OH)2}$ (white) and $\ce{Fe2O3{\cdot}xH2O}$ (brown); forms deep-blue $\ce{[Cu(NH3)4]^2+}$ and dissolves $\ce{AgCl}$ as $\ce{[Ag(NH3)2]Cl}$.
- Ostwald: $\ce{NH3 -> NO -> NO2 -> HNO3}$ (Pt/Rh, 500 K, 9 bar), NO recycled; conc. to 68%, then 98% with $\ce{H2SO4}$.
- HNO₃ as oxidiser: Cu + dilute → NO; Cu + conc. → NO₂; Zn + dilute → N₂O; Zn + conc. → NO₂. Au, Pt unaffected.
- Brown ring: $\ce{[Fe(H2O)5(NO)]^2+}$ confirms nitrate. Passivity: Fe, Cr, Al form a protective oxide film in conc. $\ce{HNO3}$.