Q.Write the balanced equation for the laboratory preparation of ammonia from ammonium chloride and slaked lime, and describe how the gas is collected and dried.

Ammonia and Nitric Acid: From Air to Fertiliser to Explosives

You already know that plants need nitrogen to grow, but the air around us — 78% nitrogen gas — is useless to them in that form. The nitrogen molecule is incredibly stable; two nitrogen atoms triple-bonded together, $N \equiv N$, simply won't react with anything mild. To make nitrogen useful, we must first break that bond and force it to combine with other elements. That is the entire point of ammonia and nitric acid production: we take inert atmospheric nitrogen and turn it into reactive chemical forms that can feed the world or, in the case of nitric acid, make explosives and fertilisers.

Ammonia by the Haber Process

The Haber process is the industrial method for fixing nitrogen — turning $N_2$ from the air into ammonia, $NH_3$. The reaction is deceptively simple:

The hydrogen comes from natural gas (methane) via steam reforming, and the nitrogen is obtained by fractional distillation of liquid air. The challenge is that this reaction is reversible and exothermic — it releases heat. Le Chatelier's principle tells us that low temperature would favour ammonia formation (since the forward reaction is exothermic), but low temperature also makes the reaction painfully slow because breaking that $N \equiv N$ triple bond requires a huge activation energy.

The solution is a compromise: a temperature around 450–500°C, a very high pressure of 150–200 atm, and an iron catalyst (with promoters like $K_2O$ and $Al_2O_3$) to speed things up. Under these conditions, about 15–20% of the reactants convert to ammonia per pass. The unreacted gases are recycled, so overall conversion is nearly 100%.

The ammonia is removed by cooling the gas mixture — ammonia liquefies at about -33°C at atmospheric pressure, while nitrogen and hydrogen remain gases — and the unreacted gases are pumped back into the reactor.

Nitric Acid by the Ostwald Process

Now take that ammonia and oxidise it. That is the Ostwald process: burning ammonia in air over a platinum-rhodium catalyst to produce nitric oxide, which then goes through a series of spontaneous reactions to become nitric acid.

The first step is the catalytic oxidation of ammonia:

This is done at about 900°C and just above atmospheric pressure. The hot gases are then cooled rapidly to prevent the NO from decomposing back. The nitric oxide then reacts with more oxygen from the air to form nitrogen dioxide:

This happens spontaneously as the gases cool. Finally, the nitrogen dioxide is absorbed in water to give nitric acid:

Notice that NO is regenerated in this last step — it gets recycled back to be oxidised again. The overall reaction from ammonia to nitric acid is:

The product is typically about 60–70% nitric acid, which can be concentrated by distillation with concentrated sulphuric acid to obtain nearly 98% nitric acid.

Why This Matters

Ammonia is the starting point for almost all nitrogen fertilisers — urea, ammonium nitrate, ammonium sulphate. Nitric acid is used to make ammonium nitrate fertiliser (by neutralising ammonia with nitric acid) and also explosives like TNT and nitroglycerin. The two processes are linked: Haber gives you ammonia, Ostwald turns some of that ammonia into nitric acid, and then you can combine them to make ammonium nitrate.

The entire modern food system depends on these two processes. Without the Haber process, the Earth could only support about 4 billion people — we are at 8 billion because we learned to pull nitrogen out of the air.

Preparation of ammonia and nitric acid is one of the most repeated industrial-chemistry questions in the NCERT Class 12 p-Block Elements chapter, searched as "ammonia and nitric acid preparation class 12 chemistry" and "Haber process Ostwald process important questions CBSE".