Reactions and uses of ammonia
Ammonia is the odd one out among the exam-syllabus gases: every other common gas you meet — carbon dioxide, sulphur dioxide, hydrogen chloride — is acidic or neutral, while ammonia is distinctly alkaline. That single fact is the key to almost everything it does. It neutralises acids to form ammonium salts, it fumes white the instant it meets hydrogen chloride gas, it burns in pure oxygen (never in ordinary air) with a pale yellowish-green flame, and — most usefully for the analytical chemistry unit — its aqueous solution precipitates the hydroxides of several common metal ions, some of which then dissolve back into solution when more ammonia is added. Dissolved in water it becomes ammonium hydroxide, the laboratory's everyday weak alkali. This article works through each reaction, ties them to ammonia's alkaline character, and closes with its major industrial and laboratory uses.
Ammonia's alkaline character (it readily accepts a proton) explains every one of its reactions — with acids, with HCl gas, on metal salt solutions, and in water — and underlies its main industrial and laboratory uses.
The core idea
Reaction with acids — forms ammonium salts. Ammonia neutralises acids exactly as a base would, giving an ammonium salt with no water produced (unlike a typical alkali-acid neutralisation, since NH₃ itself supplies no hydroxide ion until dissolved):
where: NH₃ is ammonia gas, and the product in each case is the corresponding ammonium salt — ammonium chloride or ammonium sulphate.
Reaction with HCl gas — the white fumes test. When ammonia gas meets hydrogen chloride gas directly (no solvent needed), they combine instantly to give dense white fumes of solid ammonium chloride. This is the standard mutual confirmatory test: a rod dipped in ammonia held near a rod dipped in concentrated hydrochloric acid produces visible white smoke without either gas needing to be identified first.
Combustion in oxygen. Ammonia will not sustain a flame in ordinary air, but burns readily in pure oxygen with a pale yellowish-green flame, giving nitrogen gas and water vapour:
where: the nitrogen in NH₃ is oxidised to N₂ (not to an oxide of nitrogen, since no catalyst is present), and the hydrogen is oxidised to water vapour.
Action on metal salt solutions. Ammonium hydroxide added drop by drop to solutions of metal salts such as Cu²⁺, Fe²⁺, Fe³⁺, Zn²⁺, Pb²⁺ or Ca²⁺ precipitates the corresponding metal hydroxide, because it supplies OH⁻ ions once dissolved. Adding a large excess dissolves some of these precipitates again — copper hydroxide famously redissolves to a deep blue solution — while others (like iron hydroxides) remain insoluble however much ammonia is added. This dropwise-versus-excess behaviour is the basis of several cation identification tests in the Analytical Chemistry unit.
Aqueous ammonia (ammonium hydroxide). Dissolving ammonia gas in water gives ammonium hydroxide, a weak alkali held in a reversible equilibrium:
It turns moist red litmus blue, has a sharp pungent smell, and is the laboratory's standard reagent for precipitating metal hydroxides.
Uses of ammonia. Manufacture of nitrogenous fertilisers (ammonium sulphate, urea), manufacture of nitric acid via the Ostwald process, use as a refrigerant gas, manufacture of washing and cleaning agents, and as the laboratory reagent ammonium hydroxide for precipitating metal hydroxides in qualitative analysis.
The goal is to turn the balanced equation for ammonia reacting with an acid into an actual mass calculation, using the mole concept the way an ICSE numerical question would test it — starting from a known mass of ammonia gas and finishing with the mass of the ammonium salt it produces, using only molar masses and the 1:1 mole ratio read off the balanced equation.
Step 1. Write the balanced equation and read off the mole ratio. For ammonia reacting completely with hydrogen chloride gas, , so one mole of NH₃ gives exactly one mole of NH₄Cl — a 1:1 ratio, the simplest case a mole calculation can present.
Step 2. Convert the given mass of ammonia to moles using its molar mass, (N is 14, three H is 3). For a sample of mass , moles of NH₃ .
Step 3. Apply the mole ratio from Step 1: moles of NH₄Cl produced equal moles of NH₃ consumed, since the ratio is 1:1.
Step 4. Convert moles of NH₄Cl back to mass using its molar mass, (14 + 4 + 35.5). Mass of NH₄Cl (moles of NH₄Cl) . This four-step chain — mass to moles, apply the ratio, moles to mass — is exactly the pattern examiners test for any acid-base or combustion equation in this chapter.
Worked example
Problem: 8.5 g of ammonia gas is passed into an excess of hydrogen chloride gas until the reaction is complete. Calculate the mass of ammonium chloride formed. (Take molar masses: N = 14, H = 1, Cl = 35.5.)
Step 1. Write the equation: , a 1:1 mole ratio, exactly as used in the derivation above.
Step 2. Find moles of NH₃ supplied: , so moles of NH₃ .
Step 3. Apply the 1:1 ratio: moles of NH₄Cl formed , since HCl is in excess and every mole of NH₃ is consumed.
Step 4. Convert to mass: , so mass of NH₄Cl . The reaction produces 26.75 g of ammonium chloride.
Competitive stretch
If the same 8.5 g of ammonia had instead been reacted with excess sulphuric acid using , the mole ratio changes to 2:1, so NH₃ gives only of . With , the mass formed is — always check the mole ratio from the balanced equation before converting to mass, since it is rarely 1:1.
Common traps
- Assuming ammonia burns readily in air like a fuel gas — it does not; it needs pure oxygen to sustain combustion.
- Describing the metal-salt-solution reaction as only giving a precipitate, without the dropwise-versus-excess distinction — some precipitates like copper hydroxide redissolve in excess ammonium hydroxide, others do not.
- Confusing ammonium hydroxide (NH₄OH, the aqueous solution) with ammonia gas (NH₃) when writing equations — the correct formula and state matter.
- Skipping the mole-ratio step in a stoichiometry calculation and assuming every reaction is 1:1 — check the balanced equation first, as the sulphuric acid reaction above shows.
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Why does ammonia burn in pure oxygen but not in ordinary air?
Ordinary air is only about one-fifth oxygen, heavily diluted by nitrogen and other gases, so the concentration of oxygen present is too low to sustain a spreading flame under normal conditions. Pure oxygen supplies a much higher, concentrated amount of oxidiser, which is enough to sustain and propagate the combustion reaction, giving the pale yellowish-green flame and the products nitrogen gas and water vapour.
Why does copper hydroxide redissolve in excess ammonium hydroxide but some other metal hydroxides do not?
At the syllabus level you only need to observe and describe this: a few drops of NH₄OH precipitate copper hydroxide, and a large excess dissolves that precipitate again to a deep blue solution, because copper hydroxide reacts further with the extra ammonia present. Iron hydroxides and several other metal hydroxides do not behave this way and stay as precipitates however much ammonia is added, so the redissolving behaviour becomes a distinguishing test for copper ions specifically. Writing the formula of the soluble complex ion is beyond what is required here.
Is ammonium hydroxide the same thing as ammonia?
No. Ammonia (NH₃) is the gas itself, with its own reactions such as combustion and combination with HCl gas. Ammonium hydroxide (NH₄OH) is what forms when that gas dissolves in water, , and it is this aqueous solution that behaves as a weak alkali, turns red litmus blue, and is used to precipitate metal hydroxides. Keep the two formulas and states distinct when writing equations, since examiners specifically check this.