Preparation of ammonia

The reaction. Ammonia is generated in the lab by heating a mixture of an ammonium salt and a base. The standard combination, and the one you should be able to reproduce from memory, is ammonium chloride with calcium hydroxide (slaked lime):

where: $\text{NH}_4\text{Cl}$ is ammonium chloride (the ammonium salt), $\text{Ca(OH)}_2$ is slaked lime (the base), $\Delta$ over the arrow means the mixture is gently heated, $\text{CaCl}_2$ is calcium chloride (a soluble, non-volatile by-product left behind), and $\text{NH}_3\uparrow$ is the ammonia gas that escapes and is collected. Notice the mixture must be dry, finely powdered solids — not solutions. Ammonia is exceptionally soluble in water, so if aqueous solutions were heated instead, most of the gas formed would simply redissolve in the water present before it could escape, giving a poor yield.

Two design choices, one reason. Every downstream choice in this preparation traces back to ammonia being a basic gas: Drying the gas has to avoid any acidic desiccant. Concentrated sulphuric acid and anhydrous calcium chloride are both routinely used to dry other gases, but both would react with or trap NH₃, so the correct drying agent is quicklime (CaO), a basic solid that removes moisture without consuming the gas. Collecting the gas has to account for density and solubility together. NH₃ is markedly lighter than air and dissolves readily in water, so it is collected by upward displacement of air: the delivery tube feeds gas into a jar held mouth-downward, letting the light gas rise and fill the jar from the top while denser air is pushed out below. Collecting it over a water trough, the method used for less-soluble gases like hydrogen, would waste almost all the gas as it dissolved.

Every choice in this preparation is a direct consequence of two physical facts about ammonia: it is a base, and it is far lighter and far more soluble than air's other constituents.

Why the reactants are dry solids. Ammonia dissolving readily in water is the same property that makes the classic "fountain experiment" work — but here that solubility is a liability, not a feature. Any water present, whether from an aqueous solution of the reactants or from moisture in the apparatus, would absorb newly formed NH₃ before it could reach the collecting jar. Using dry, finely powdered, thoroughly mixed solids ensures the ammonia escapes as gas the instant it forms.

Why slaked lime and not another base. Ca(OH)₂ is cheap, safe to handle, and reacts with ammonium salts on gentle heating without side reactions. Its solid by-product, calcium chloride, is not volatile, so it stays behind in the reaction tube rather than contaminating the gas stream — the ammonia that reaches the drying tube is essentially pure.

Why quicklime and not an acidic desiccant. A drying agent's only job is to remove water vapour without touching the gas itself. Concentrated sulphuric acid and anhydrous calcium chloride do that job for acidic or neutral gases, but both are (or behave as) acids toward ammonia, so they would react with or bind the very gas being dried, sharply cutting the yield. Quicklime is basic, so it is chemically inert toward NH₃ while still soaking up moisture as it slakes to Ca(OH)₂.

Why upward displacement of air, not water. Two properties decide the collection method for any gas: is it denser or lighter than air, and how soluble is it in water. Ammonia is both markedly lighter than air and extremely soluble in water, so the only sensible method is to trap the rising gas in an inverted jar held mouth-down, letting air escape from below as the lighter NH₃ accumulates at the top.

Problem. A student is asked to prepare and identify ammonia gas in the laboratory. Write the balanced equation for the preparation, name the correct drying agent with a reason, state the collection method with a reason, and describe one confirmatory test for the gas collected.

Step 1. Identify the reactants: ammonium chloride (the ammonium salt) and calcium hydroxide, i.e. slaked lime (the base), used as dry powdered solids and mixed thoroughly before heating.

Step 2. Write the balanced equation for the reaction on gentle heating:

The coefficient 2 on both NH₄Cl and NH₃ is essential — losing it is the single most common slip in this equation.

Step 3. Choose the drying agent. Since ammonia is a basic gas, it cannot be dried with concentrated sulphuric acid, because that would consume the gas by the reaction $2\text{NH}_3 + \text{H}_2\text{SO}_4 \rightarrow (\text{NH}_4)_2\text{SO}_4$. The correct drying agent is quicklime (CaO), which removes moisture without reacting with NH₃.

Step 4. Choose the collection method. Because ammonia is lighter than air and extremely soluble in water, it is collected by upward displacement of air, with the gas jar held mouth-downward, and never over a water trough.

Step 5. State a confirmatory test. Bring a glass rod dipped in concentrated hydrochloric acid near the mouth of the gas jar; dense white fumes of ammonium chloride form immediately: $\text{NH}_3 + \text{HCl} \rightarrow \text{NH}_4\text{Cl (white fumes)}$. Alternatively, moist red litmus paper held near the gas turns blue, confirming an alkaline gas.

The prepared gas is confirmed as ammonia by the white-fumes test (or the litmus test), obtained via the equation $2\text{NH}_4\text{Cl} + \text{Ca(OH)}_2 \xrightarrow{\Delta} \text{CaCl}_2 + 2\text{NH}_3\uparrow + 2\text{H}_2\text{O}$, dried over quicklime, and collected by upward displacement of air.

Competitive stretch

Problem. Why would a student who prepares ammonia using aqueous solutions of ammonium chloride and calcium hydroxide, instead of dry solids, obtain a much smaller volume of gas even though the same balanced equation applies?

Step 1. The balanced equation is unchanged — the stoichiometry of moles of NH₃ produced per mole of NH₄Cl reacted stays $2:2$, i.e. one mole of NH₃ per mole of NH₄Cl.

Step 2. The difference lies entirely in solubility, not stoichiometry. Ammonia is extremely soluble in water. When the reaction happens in aqueous solution, most of the NH₃ formed dissolves straight back into the surrounding water instead of escaping as gas.

Step 3. Only the small fraction of NH₃ that reaches saturation and escapes the liquid surface is available to collect, so the gas volume collected is far below what the equation predicts. The lower yield is a solubility effect, not a change in the chemistry — using dry solids avoids it entirely.

• Using concentrated sulphuric acid or anhydrous calcium chloride to dry ammonia — both react with or absorb NH₃; the correct drying agent is quicklime (CaO).
• Collecting ammonia the way HCl gas is collected (downward delivery) — wrong, because ammonia is lighter than air, so it needs upward displacement of air with the jar mouth held down.
• Dropping the coefficient 2 in front of NH₄Cl and NH₃ when writing the equation — a simple but heavily penalised balancing slip.
• Mixing up this laboratory preparation (ammonium salt + base, small scale) with the industrial Haber process (N₂ + H₂ under pressure and a catalyst) — they are different methods for different purposes.
• Trying the preparation with aqueous solutions instead of dry solids — this lowers the yield because so much of the very soluble ammonia redissolves before it can be collected as gas.

Ask the tutor: if the "basic gas needs a basic drying agent, never an acidic one" logic still feels shaky, ask the AI tutor to walk through two or three more drying-agent and collection-method drills for other gases.