Nitric Acid Dissolves Copper Without Releasing Any Hydrogen

Drop a copper coin into dilute hydrochloric acid and nothing happens — copper lies below hydrogen and cannot displace it. Drop the same coin into nitric acid and it dissolves, the solution turns blue and a brown or colourless gas bubbles off. Yet none of that gas is hydrogen.

Nitric acid is two things at once.

- It is an acid, giving hydronium ions in water and reacting with bases, oxides and carbonates like any other acid
- It is a powerful oxidising agent, able to supply oxygen to metals and non-metals

The second property explains the copper coin. Nitric acid oxidises the copper directly, and any hydrogen that might have formed is oxidised to water, so hydrogen gas is almost never seen. The gas that does appear is an oxide of nitrogen, formed as nitric acid itself is reduced.

The same oxidising power shapes how nitric acid is handled.

- It attacks rubber and cork, so it must be prepared in an all-glass apparatus
- It slowly decomposes, especially in light, which is why it turns yellow and is stored in dark bottles

Nitric acid is made on a huge scale, mostly by Ostwald's process, which starts from the ammonia of the previous parts. Its largest use is making fertilisers such as ammonium nitrate; it is also used for explosives, dyes and in metal work.

This part covers:

- Laboratory preparation from potassium or sodium nitrate
- Ostwald's process — the conversion of ammonia into nitric acid
- Nitric acid as an oxidising agent, with copper, carbon and sulphur
- Its acidic properties, and the tests for nitric acid and nitrate ions

The link that runs through the chapter. Hydrogen chloride came from salt and sulphuric acid; nitric acid comes from a nitrate and sulphuric acid by the same principle. Ammonia's catalytic oxidation from Part 3 becomes the first step of Ostwald's process. The four compounds are one connected story.

This page covers the ICSE Class 10 Chemistry study of nitric acid: laboratory preparation, Ostwald's process, oxidising properties, acidic properties and tests.

Potassium nitrate or sodium nitrate is heated with concentrated sulphuric acid below about 200 °C in an all-glass retort; nitric acid vapour distils over and is condensed in a cooled receiver as a yellowish liquid.

Reactants: potassium nitrate or sodium nitrate, and concentrated sulphuric acid.

Condition: heating, keeping the temperature **below about $200\text{ °C}$**:

$${\text{KNO}}_3+{\text{H}}_2{\text{SO}}_4\to {\text{KHSO}}_4+{\text{HNO}}_3$$
$${\text{NaNO}}_3+{\text{H}}_2{\text{SO}}_4\to {\text{NaHSO}}_4+{\text{HNO}}_3$$

Why it works. Concentrated sulphuric acid is far less volatile than nitric acid, so on heating the more volatile nitric acid distils out — the same principle as the preparation of hydrogen chloride from salt.

Apparatus:

- A glass retort with a glass stopper, containing the nitrate and acid
- A receiver flask into which the neck of the retort leads
- The receiver is cooled under running water or with ice, so that nitric acid vapour condenses to a liquid

Why all glass. Nitric acid vapour attacks rubber and cork, oxidising them. Stoppers and joints must therefore be glass, and a retort with a long glass neck avoids the need for any other connection.

Why the temperature is kept low:

- Nitric acid decomposes at higher temperatures, so less would be collected
- Stronger heating wastes fuel and risks cracking the glass retort

Collection: as a liquid in the cooled receiver.

Why the acid looks yellow. Some nitric acid decomposes during the preparation, and the reddish-brown nitrogen dioxide dissolves in the acid, colouring it:

$$4{\text{HNO}}_3\to 4{\text{NO}}_2+2{\text{H}}_2\text{O}+{\text{O}}_2$$

The yellow colour can be removed by bubbling dry air or carbon dioxide through the warm acid, which sweeps out the dissolved nitrogen dioxide.

Precautions:

- **Temperature below about $200\text{ °C}$
- All-glass apparatus, with no rubber or cork
- Receiver kept cold
- Nitric acid is highly corrosive and must not touch the skin

Identification: warm the liquid with copper turnings — reddish-brown fumes of nitrogen dioxide and a blue-green solution show it is nitric acid.

Worked example — yield.** What mass of nitric acid can be obtained from $20.2\text{ g}$ of potassium nitrate? K $39$, N $14$, O $16$, H $1$.

$$M({\text{KNO}}_3)=39+14+48=101\text{moles}=\frac{20.2}{101}=0.2$$
$${\text{HNO}}_3=0.2\text{ mol}=0.2\times 63=12.6\text{ g}$$

Worked check — balancing the decomposition equation. Hydrogen $4=4$; nitrogen $4=4$; oxygen $12$ on the left and $8+2+2=12$ on the right. Balanced.

An everyday example. Bottles of concentrated nitric acid in a school laboratory are brown or dark glass, and older bottles often show a yellowish liquid. Light speeds up the decomposition above, so the dark glass slows the yellowing.

The boundary case. Nitric acid made this way is concentrated, but not pure: it contains dissolved nitrogen dioxide. The yellow colour is an impurity, not a property of pure nitric acid, which is colourless.

Ammonia is oxidised by air over platinum gauze at about 800 °C to nitric oxide, which combines with more oxygen to nitrogen dioxide, and the nitrogen dioxide is absorbed in water with air to give nitric acid.

Raw materials: ammonia — from Haber's process — air, and water.

Step 1 — catalytic oxidation of ammonia. Dry, purified ammonia and excess air are passed through a converter over platinum gauze heated to about $800\text{ °C}$:

$$4{\text{NH}}_3+5{\text{O}}_2\to 4\text{NO}+6{\text{H}}_2\text{O}+\text{heat}$$

The reaction is exothermic, so once started it keeps the gauze hot without further external heating.

Step 2 — oxidation of nitric oxide. The hot gases are cooled, and the nitric oxide combines with more oxygen from the excess air:

$$2\text{NO}+{\text{O}}_2\to 2{\text{NO}}_2$$

Step 3 — absorption in water. The nitrogen dioxide, with air, rises through an absorption tower packed with quartz pieces, over which water trickles down:

$$4{\text{NO}}_2+2{\text{H}}_2\text{O}+{\text{O}}_2\to 4{\text{HNO}}_3$$

The product is dilute nitric acid, which can be concentrated by distillation.

The plant, as a flow of stages:

- Air and ammonia mixed and purified
- Converter with platinum gauze — nitric oxide formed
- Cooler and oxidation chamber — nitrogen dioxide formed
- Absorption tower with quartz packing and water — nitric acid formed
- Concentration by distillation

Why the quartz packing. It gives a large surface area over which the gas and the trickling water meet, so absorption is efficient, and quartz is not attacked by the acid.

Worked check — the overall change. Adding the three steps, with step 2 doubled to match:

$$4{\text{NH}}_3+5{\text{O}}_2\to 4\text{NO}+6{\text{H}}_2\text{O}$$
$$4\text{NO}+2{\text{O}}_2\to 4{\text{NO}}_2$$
$$4{\text{NO}}_2+2{\text{H}}_2\text{O}+{\text{O}}_2\to 4{\text{HNO}}_3$$

Cancelling the intermediates and the water on both sides:

$$4{\text{NH}}_3+8{\text{O}}_2\to 4{\text{HNO}}_3+4{\text{H}}_2\text{O}⟹{\text{NH}}_3+2{\text{O}}_2\to {\text{HNO}}_3+{\text{H}}_2\text{O}$$

Worked example — nitric acid from ammonia. What mass of nitric acid can be made from $17\text{ kg}$ of ammonia, if conversion is complete?

$$17{\text{ kg NH}}_3=1\text{ kmol}{\text{HNO}}_3=1\text{ kmol}=63\text{ kg}$$

Worked check — balancing step 3. Nitrogen $4=4$; oxygen $8+2+2=12$ on the left and $4\times 3=12$ on the right; hydrogen $4=4$. Balanced.

An everyday connection. Much of this nitric acid is neutralised with ammonia to make ammonium nitrate fertiliser — so a fertiliser plant may run Haber's process and Ostwald's process side by side, turning air, water and natural gas into a nitrogen fertiliser.

The boundary case. The platinum catalyst decides the product of step 1. Without it, ammonia would burn to nitrogen and water rather than to nitric oxide, and the nitrogen could never be turned into nitric acid — the catalyst makes the whole process possible, not just faster.

Nitric acid readily gives up oxygen, forming nitrogen dioxide when concentrated and nitric oxide when dilute; it oxidises copper to copper nitrate, carbon to carbon dioxide and sulphur to sulphuric acid.

The source of the oxidising power. Nitric acid decomposes to give nascent oxygen, a very reactive form:

Concentrated acid:

$$2{\text{HNO}}_3\to 2{\text{NO}}_2+{\text{H}}_2\text{O}+[\text{O}]$$

Dilute acid:

$$2{\text{HNO}}_3\to 2\text{NO}+{\text{H}}_2\text{O}+3[\text{O}]$$

1. With copper — concentrated acid.

$$\text{Cu}+4{\text{HNO}}_3\to {\text{Cu(NO}}_3{)}_2+2{\text{H}}_2\text{O}+2{\text{NO}}_2$$

Observations: copper dissolves, reddish-brown fumes of nitrogen dioxide, and a blue-green solution.

2. With copper — dilute acid.

$$3\text{Cu}+8{\text{HNO}}_3\to 3{\text{Cu(NO}}_3{)}_2+4{\text{H}}_2\text{O}+2\text{NO}$$

Observations: copper dissolves slowly, a blue solution forms, and a colourless gas, nitric oxide, is released — which turns reddish-brown at the mouth of the tube as it meets air: $2\text{NO}+{\text{O}}_2\to 2{\text{NO}}_2$.

3. With carbon — hot concentrated acid.

$$\text{C}+4{\text{HNO}}_3\to {\text{CO}}_2+2{\text{H}}_2\text{O}+4{\text{NO}}_2$$

Carbon is oxidised to carbon dioxide, with brown fumes of nitrogen dioxide.

4. With sulphur — hot concentrated acid.

$$\text{S}+6{\text{HNO}}_3\to {\text{H}}_2{\text{SO}}_4+2{\text{H}}_2\text{O}+6{\text{NO}}_2$$

Sulphur is oxidised to sulphuric acid, with brown fumes.

Worked check — balancing the dilute copper equation.

- Copper: $3=3$
- Hydrogen: $8$ on the left; $4\times 2=8$ on the right
- Nitrogen: $8$ on the left; $3\times 2+2=8$ on the right
- Oxygen: $8\times 3=24$ on the left; $3\times 6+4+2=24$ on the right

Balanced.

Worked example — gas from copper. What volume of nitric oxide at STP is released when $19.05\text{ g}$ of copper dissolves completely in dilute nitric acid? Cu $63.5$.

$$\frac{19.05}{63.5}=0.3\text{ mol Cu}\text{NO}=0.3\times \frac{2}{3}=0.2\text{ mol}=0.2\times 22.4=4.48\text{ L}$$

Why hydrogen is not released. With most metals, any hydrogen formed is immediately oxidised to water by the nascent oxygen. So nitric acid gives oxides of nitrogen, not hydrogen, even with metals above hydrogen in the activity series.

An everyday example. Artists and printmakers etch designs into copper plates with nitric acid: the acid dissolves the exposed copper, while areas protected by a wax coating stay untouched, leaving the design cut into the metal.

The boundary case — passivity. Concentrated nitric acid does not dissolve iron or aluminium as expected; it forms a thin, protective oxide layer that stops further attack. That is why concentrated nitric acid can be transported in aluminium containers, even though dilute nitric acid attacks aluminium.

Very dilute nitric acid gives hydrogen with magnesium and manganese, and dilute nitric acid neutralises oxides and hydroxides and releases carbon dioxide from carbonates and sulphur dioxide from sulphites; nitrate ions are identified by the brown ring test.

1. With magnesium and manganese — the exception that gives hydrogen. With very dilute nitric acid in the cold, these two metals release hydrogen:

$$\text{Mg}+2{\text{HNO}}_3\to {\text{Mg(NO}}_3{)}_2+{\text{H}}_2$$
$$\text{Mn}+2{\text{HNO}}_3\to {\text{Mn(NO}}_3{)}_2+{\text{H}}_2$$

2. With oxides and hydroxides — salt and water.

$$\text{CuO}+2{\text{HNO}}_3\to {\text{Cu(NO}}_3{)}_2+{\text{H}}_2\text{O}$$
$$\text{ZnO}+2{\text{HNO}}_3\to {\text{Zn(NO}}_3{)}_2+{\text{H}}_2\text{O}$$
$$\text{NaOH}+{\text{HNO}}_3\to {\text{NaNO}}_3+{\text{H}}_2\text{O}$$
$${\text{Ca(OH)}}_2+2{\text{HNO}}_3\to {\text{Ca(NO}}_3{)}_2+2{\text{H}}_2\text{O}$$

3. With carbonates and hydrogen carbonates — salt, water and carbon dioxide.

$${\text{Na}}_2{\text{CO}}_3+2{\text{HNO}}_3\to 2{\text{NaNO}}_3+{\text{H}}_2\text{O}+{\text{CO}}_2$$
$${\text{CaCO}}_3+2{\text{HNO}}_3\to {\text{Ca(NO}}_3{)}_2+{\text{H}}_2\text{O}+{\text{CO}}_2$$
$${\text{NaHCO}}_3+{\text{HNO}}_3\to {\text{NaNO}}_3+{\text{H}}_2\text{O}+{\text{CO}}_2$$

4. With sulphites and hydrogen sulphites — salt, water and sulphur dioxide.

$${\text{Na}}_2{\text{SO}}_3+2{\text{HNO}}_3\to 2{\text{NaNO}}_3+{\text{H}}_2\text{O}+{\text{SO}}_2$$
$${\text{NaHSO}}_3+{\text{HNO}}_3\to {\text{NaNO}}_3+{\text{H}}_2\text{O}+{\text{SO}}_2$$

5. Tests for nitric acid.

- Heating concentrated nitric acid gives reddish-brown fumes of nitrogen dioxide
- Adding copper turnings to concentrated nitric acid gives reddish-brown fumes and a blue-green solution

6. Test for the nitrate ion — the brown ring test.

- To the solution of the nitrate, add a freshly prepared solution of iron(II) sulphate
- Pour concentrated sulphuric acid slowly down the side of the tilted test tube, so that it forms a separate layer at the bottom
- A brown ring forms where the two layers meet, confirming the nitrate ion

The chemistry of the brown ring. Sulphuric acid releases nitric acid from the nitrate; iron(II) sulphate reduces it to nitric oxide; the nitric oxide combines with more iron(II) sulphate to form the brown compound:

$$2{\text{NaNO}}_3+{\text{H}}_2{\text{SO}}_4\to {\text{Na}}_2{\text{SO}}_4+2{\text{HNO}}_3$$
$$6{\text{FeSO}}_4+3{\text{H}}_2{\text{SO}}_4+2{\text{HNO}}_3\to 3{\text{Fe}}_2({\text{SO}}_4{)}_3+4{\text{H}}_2\text{O}+2\text{NO}$$
$${\text{FeSO}}_4+\text{NO}\to {\text{FeSO}}_4\cdot \text{NO}$$

Worked check — balancing the iron(II) sulphate equation.

- Iron: $6$ on the left; $3\times 2=6$ on the right
- Sulphur: $6+3=9$ on the left; $3\times 3=9$ on the right
- Hydrogen: $6+2=8$ on the left; $4\times 2=8$ on the right
- Nitrogen: $2=2$
- Oxygen: $24+12+6=42$ on the left; $36+4+2=42$ on the right

Balanced.

Worked example — gas from a carbonate. What volume of carbon dioxide at STP is released when $5.3\text{ g}$ of sodium carbonate reacts with excess dilute nitric acid, and what mass of nitric acid is used? Na $23$, C $12$.

$$M({\text{Na}}_2{\text{CO}}_3)=106\frac{5.3}{106}=0.05\text{ mol}$$
$${\text{CO}}_2=0.05\times 22.4=1.12\text{ L}{\text{HNO}}_3=0.1\text{ mol}=6.3\text{ g}$$

An everyday example. Soil and water laboratories test for nitrate because excess fertiliser can wash nitrate into wells and ponds. The brown ring test is the classic chemical way to detect it, though modern laboratories also use instruments.

The boundary cases. Only magnesium and manganese, with very dilute acid, give hydrogen — all other metals give oxides of nitrogen. And with sulphites, concentrated nitric acid can oxidise the sulphur dioxide further, so these acidic reactions are described with dilute acid.