Conversion to oxide

An ore fresh from concentration is still, chemically, in the wrong form. Metals are almost always pulled out of their compounds by reduction of an oxide — but real ores turn up as carbonates, sulphides, or hydrated oxides, never conveniently as pure oxide on their own. Somebody has to bridge that gap, and the bridge is heat. Two heating processes do the job: calcination, a gentle bake in limited air that drives off carbon dioxide or water, and roasting, a fierce bake in excess air that burns sulphur away as sulphur dioxide. Both leave the metal behind as an oxide, ready for the next stage of metallurgy. Mix them up under exam pressure and you lose easy marks — the ore type and the air supply are the two facts that decide everything here.

The core insight

Carbonate and hydrated-oxide ores are calcined in limited air to drive off CO₂ or water vapour, while sulphide ores are roasted in excess air to burn off sulphur as SO₂ — both processes converting the concentrated ore into the metal oxide needed for reduction.

The core idea

Why an oxide at all? The reduction step that follows concentration is built to strip oxygen from a metal oxide, not to strip carbonate or sulphide groups. So before reduction can start, whatever anion is attached to the metal — carbonate, hydrated oxide, or sulphide — has to be swapped out for plain oxide. That swap happens by heating, and which of the two processes is used depends entirely on what kind of ore you are holding.

Calcination — limited air, carbonates and hydrated oxides. The concentrated ore is heated strongly with little or no air, usually in a reverberatory furnace. There is no sulphur to burn, so no oxygen is needed for combustion; heat alone is enough to drive off the volatile part.

Hydrated oxide (bauxite, a hydrated oxide of aluminium):

\text{Al}_2\text{O}_3\cdot 2\text{H}_2\text{O} \xrightarrow{\Delta} \text{Al}_2\text{O}_3 + 2\text{H}_2\text{O}\uparrow

where: the arrow with $\Delta$ marks strong heating, and the water leaves as steam, leaving anhydrous alumina behind.

Carbonate (zinc carbonate, calamine):

\text{ZnCO}_3 \xrightarrow{\Delta} \text{ZnO} + \text{CO}_2\uparrow

where: carbon dioxide gas escapes and the solid residue is zinc oxide, the oxide needed for reduction.

Roasting — excess air, sulphides. Here the concentrated ore is heated strongly with a plentiful supply of air, because oxygen must actually combine with the ore for the reaction to happen — the sulphide is oxidised.

2\text{ZnS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{ZnO} + 2\text{SO}_2\uparrow

where: oxygen from the excess air combines with both the zinc and the sulphur; zinc ends up bonded to oxygen (the desired oxide) while sulphur escapes as sulphur dioxide gas. The single feature that separates the two processes is the air supply — calcination avoids it, roasting depends on it — and that single difference follows directly from what has to leave the solid: a gas that was already loosely attached (CO₂, H₂O) needs only heat, while sulphur bonded to the metal needs oxygen to burn it away.

Figure

Two furnace diagrams side by side, calcination on the left showing a carbonate ore heated with a small air symbol releasing carbon dioxide gas, and roasting on the right showing a sulphide ore heated with a large air symbol releasing sulphur dioxide gas, both ending in a metal oxide

Why it works

There is no algebraic formula to derive here — this is a descriptive step in the metallurgy sequence, and the reasoning that makes it click is chemical, not mathematical.

Step 1. Start from the destination, not the start. Reduction — the step right after this one — is built around removing oxygen from a metal oxide, typically using carbon or a more reactive metal. That machinery only works if the metal is already bonded to oxygen. So whatever the ore's natural form, it has to be converted to oxide first; this stage exists purely to make the next stage possible.

Step 2. Sort the ore by what it is bonded to. A carbonate ore has the metal bonded to a carbonate group (CO₃²⁻); a hydrated oxide has the metal already bonded to oxide, just with extra water molecules attached; a sulphide ore has the metal bonded to sulphur. Each of these three "wrong forms" needs a different push to become plain oxide.

Step 3. Match the push to the bond. Carbonate and water are both weakly held and volatile — heat alone is enough to shake them loose and drive them off as gas, with no chemical reaction against oxygen needed. That is calcination, and it is deliberately done in limited air so nothing extra combines with the ore. Sulphur, by contrast, is not volatile on its own; it has to be chemically converted into a gas by combining it with oxygen to form SO₂. That combination is a combustion reaction, so it requires exactly what calcination avoids — a generous, excess supply of air. That is roasting.

Step 4. Read off the outcome. In both cases the metal itself stays behind in the solid, now bonded only to oxygen — the oxide. The gas that leaves (CO₂, H₂O vapour, or SO₂) is simply the anion the metal was originally attached to, driven off in whichever way its chemistry allows. Once you see that the two processes are just two different answers to "how do I remove this particular anion," the choice between calcination and roasting stops being a fact to memorise and becomes a fact you can work out from the ore's formula.

Worked example

Question: Zinc blende (ZnS), after concentration by froth flotation, must be converted to zinc oxide before reduction. Explain which process is used, describe the conditions, and write the balanced equation.

Step 1. Zinc blende is a sulphide ore, so the appropriate process is roasting, not calcination. Calcination is reserved for carbonate and hydrated-oxide ores and involves no combustion with oxygen, so it cannot remove the sulphur bonded to the zinc here.

Step 2. In roasting, the concentrated ZnS is heated strongly in a plentiful supply of air (oxygen), typically in a furnace designed to allow good air circulation over the ore.

Step 3. The oxygen reacts with the sulphide, oxidising the zinc to zinc oxide while the sulphur is oxidised and escapes as sulphur dioxide gas:

2\text{ZnS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{ZnO} + 2\text{SO}_2\uparrow

Step 4. The zinc oxide produced is now in the correct chemical form to be reduced (commonly using carbon) to give free zinc metal in the next metallurgical step. Answer: zinc blende is roasted in excess air to give ZnO, with SO₂ escaping as the by-product gas.

Competitive stretch

If instead the ore given were bauxite (Al₂O₃·2H₂O), the same question format flips the answer entirely: bauxite is a hydrated oxide, not a sulphide, so it is calcined, not roasted, releasing water vapour rather than SO₂, by the equation $\text{Al}_2\text{O}_3\cdot 2\text{H}_2\text{O} \xrightarrow{\Delta} \text{Al}_2\text{O}_3 + 2\text{H}_2\text{O}\uparrow$ . Examiners like to swap the ore type between otherwise identical questions precisely to check you are reasoning from the ore's formula, not reciting a memorised paragraph.

Common traps

Swapping calcination and roasting. Remember "Calcination = Carbonate, limited air"; "Roasting = suphide, excess air, R for the Reeking rotten-egg sulphur smell of escaping SO₂."
Thinking both processes need excess air. Only roasting requires excess air; calcination is done with limited or no air specifically to avoid unwanted combustion.
Forgetting bauxite is a hydrated oxide, not a carbonate. Its calcination removes water of crystallisation, not CO₂ — do not write CO₂ as the gas released from bauxite.
Missing the "why" of the step. Memorising the two equations without being able to say that the entire point is to obtain the metal in oxide form, ready for reduction, loses reasoning marks even when the equation is correct.

Ask the tutor: Give me an unfamiliar ore's formula and ask the AI tutor to identify whether it needs calcination or roasting, then check your reasoning.

How do I quickly tell which process an unfamiliar ore needs?

Look only at the anion attached to the metal in the ore's formula. If you see a carbonate group (CO₃) or the ore is described as a hydrated oxide (extra water molecules attached to an oxide), calcine it — heat in limited air is enough. If you see sulphur bonded to the metal (a sulphide, often written as MS or M₂S), roast it — you need excess air so oxygen can burn the sulphur away as SO₂. The formula tells you everything; you never need to guess.

Why does calcination avoid air if heating is happening anyway?

Because there is nothing in a carbonate or hydrated-oxide ore that needs to react with oxygen — the water or CO₂ is already loosely held and simply needs enough heat to break free as a gas. Letting extra air in during calcination is not exactly dangerous, but it adds nothing useful; it is limited or absent purely because roasting's excess-air step is not required here. The two processes are matched to two different chemical needs, not to a single shared temperature effect.

Does the ore end up as pure metal after calcination or roasting?

No — both processes stop at the oxide, not the free metal. Calcination and roasting only change the anion attached to the metal (from carbonate, hydrated oxide, or sulphide, to plain oxide); the metal itself is never reduced to its elemental form in this step. Getting the actual metal out requires a further step, reduction of the oxide (commonly with carbon or a more reactive metal), which is covered as the next stage of the metallurgy sequence.