Metals and Non-metals

Pick up a one-rupee coin, a steel spoon, a pencil's graphite core and a pinch of table salt. Every one of these objects owes its identity to whether its atoms are metal or non-metal, and to how differently those two families behave. The coin and the spoon shine, bend without snapping, and carry heat and current easily. The graphite is a strange exception — a non-metal that conducts. The salt is neither shiny nor bendable, yet it dissolves readily and conducts electricity the moment it is in solution. This chapter is the story of why. We start where Class IX left off: sorting elements by touch, sound and shine, and quickly discover that physical tests alone are unreliable, since mercury is a liquid metal, iodine is a lustrous non-metal, and graphite conducts like one. The dependable dividing line is chemical behaviour: how an element reacts with oxygen, water, acids and other metals' salts. Running those four tests produces the reactivity series, the single most useful tool in this chapter — it predicts which metal displaces which, which oxide carbon can reduce and which needs electrolysis, and why sodium lives in kerosene while gold survives centuries underground. Underneath every reaction is a simple electron story. Metal atoms give away outer electrons to reach a stable, noble-gas-like configuration; non-metal atoms accept them. That transfer builds charged ions that lock into ionic compounds — brittle, high-melting solids that conduct only once melted or dissolved, never as solids. The same reactivity series then tells metallurgists how to win a metal back from its ore, and explains why iron rusts, why aluminium does not, and why we alloy nearly every metal we actually use. By the end you should be able to read the reactivity series like a map and use it to predict a reaction, explain a bond, or choose an extraction method.

Key takeaways

Metals are lustrous, malleable, ductile, sonorous and conduct heat and electricity well; non-metals are mostly the opposite, with exceptions like mercury (liquid metal), iodine (lustrous non-metal) and graphite (conducts).
Metals react with oxygen to form mostly basic oxides; aluminium oxide and zinc oxide are amphoteric, reacting with both acids and bases.
Reactivity with water falls from potassium and sodium (violent, even cold) through calcium and magnesium (hot water or steam) to iron, aluminium and zinc (steam only) to copper, silver and gold (no reaction).
Metals above hydrogen in the reactivity series liberate hydrogen from dilute acids; copper, silver and gold do not.
A more reactive metal displaces a less reactive one from its salt solution — combined with the oxygen, water and acid tests, this builds the reactivity series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, (H), Cu, Hg, Ag, Au.
Metals form ionic compounds with non-metals by losing electrons to a stable octet while the non-metal gains them; the resulting ions attract strongly, giving high melting points and no separate molecules.
Ionic compounds are hard but brittle, generally soluble in water, and conduct only when molten or dissolved, never as solids.
A metal's series position decides its extraction: unreactive metals occur free, moderately reactive ones need carbon reduction after roasting or calcination, and the most reactive need electrolytic reduction.
Corrosion (such as rusting) needs air and moisture together, and slows with painting, oiling, galvanising, anodising or alloying.
Key idea: a metal's place in the reactivity series — set by its reactions with oxygen, water, acids and other metals' salts — governs the ionic bonds it forms, the extraction method its ore needs, and how readily it corrodes.

The core insight

The reactivity series, built from a metal's reactions with oxygen, water, acids and displacement, is the single ranking that explains how metals bond with non-metals by electron transfer, how each metal must be extracted from its ore, and how vulnerable it is to corrosion.

What this chapter covers

The chapter builds in a deliberate order. First come the physical properties that sort elements by touch, sound and shine — lustre, hardness, malleability, ductility, conduction, sonority — and the exceptions that show these tests are not enough alone. Next come the chemical properties: reactions with oxygen, water, dilute acids, and each other's salt solutions, whose combined evidence produces the reactivity series that anchors everything after it. The chapter then explains why these reactions happen, in terms of electron transfer and ionic bonding, with the resulting properties of ionic compounds. It moves on to where metals are found in nature — minerals, ores and gangue — and how the reactivity series dictates the extraction method: native occurrence, carbon reduction, or electrolysis, followed by refining. It closes with corrosion and the everyday defences against it, including the useful idea of an alloy.

The core ideas

Physical properties: metals versus non-metals

Metals, in their pure state, have a shining surface — metallic lustre — and are generally hard, though hardness varies from metal to metal. Most can be hammered into thin sheets (malleability) and drawn into thin wires (ductility); gold is the most malleable and ductile metal known, drawable from one gram into about two kilometres of wire. Metals conduct heat and electricity well (silver and copper best; lead and mercury poorly), and most ring when struck — sonorous, which is why bells are cast from metal, not sulphur. Non-metals are mostly brittle solids or gases (bromine is the one liquid non-metal), lack lustre, and conduct poorly. But the physical test alone is unreliable: mercury is a liquid metal, gallium and caesium melt on your palm, iodine is a lustrous non-metal, graphite conducts like a metal, and the alkali metals lithium, sodium and potassium are soft enough to cut with a knife. This is why the chapter turns next to chemical behaviour, which classifies elements far more reliably.

Chemical properties: reactions with oxygen, water and acids

Almost every metal combines with oxygen to form a metal oxide, and almost every metal oxide is basic.

\text{Metal} + \text{Oxygen} \xrightarrow{} \text{Metal oxide}

Copper heated in air forms black copper(II) oxide, and aluminium forms aluminium oxide:

2\text{Cu} + \text{O}_2 \xrightarrow{\Delta} 2\text{CuO}, \qquad 4\text{Al} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{Al}_2\text{O}_3

where: $\Delta$ marks heat, and (s), (l), (g), (aq) mark solid, liquid, gas and aqueous states. Aluminium oxide and zinc oxide react with both acids and bases — amphoteric oxides. Sodium and potassium oxide dissolve directly in water to form alkalis, $\text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \to 2\text{NaOH}(aq)$ . Metals also react with water to give a metal oxide or hydroxide plus hydrogen gas, not all at the same rate.

\text{Metal} + \text{Water} \to \text{Metal oxide/hydroxide} + \text{Hydrogen}

Potassium and sodium react so violently with cold water that the hydrogen catches fire; calcium reacts gently and floats as bubbles cling to it; magnesium needs hot water; aluminium, iron and zinc need steam; lead, copper, silver and gold do not react with water at all. Reactive metals also displace hydrogen from dilute acids.

\text{Metal} + \text{Dilute acid} \to \text{Salt} + \text{Hydrogen}

In dilute hydrochloric acid the bubbling rate — and so the reactivity — falls as Mg is greater than Al is greater than Zn is greater than Fe, while copper does not react. (Nitric acid usually consumes the hydrogen instead of releasing it, being a strong oxidiser, except very dilute acid with magnesium or manganese.)

The reactivity series

None of the three tests above works on every metal — copper does not react with water, silver does not react with dilute acid — so a fourth test closes the gaps: a displacement reaction, where a more reactive metal pushes a less reactive one out of its salt solution.

\text{Metal A} + \text{Salt solution of B} \to \text{Salt solution of A} + \text{Metal B}

If A displaces B, then A is more reactive than B. Combining the oxygen, water, acid and displacement evidence gives the standard reactivity (activity) series, most to least reactive: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, (hydrogen), copper, mercury, silver, gold — the tool used later to decide each metal's extraction method.

Figure

The reactivity series from potassium to gold shown as a vertical ladder, divided into three coloured zones for electrolytic extraction, extraction by reduction with carbon, and metals found free in nature

Ionic bonding: why metals and non-metals react

The tendency to react is a tendency to attain a stable, completely filled outermost shell like a noble gas. Sodium has only one electron in its outer shell; losing it leaves a full shell underneath and a net positive charge, the sodium cation $\text{Na}^+$ . Chlorine has seven outer electrons; gaining one completes its octet and leaves a net negative charge, the chloride anion $\text{Cl}^-$ .

\text{Na} \to \text{Na}^+ + e^-, \qquad \text{Cl} + e^- \to \text{Cl}^-

where: $e^-$ is one electron transferred directly from metal to non-metal — none is created or destroyed, only relocated. The oppositely charged ions then attract by strong electrostatic force to form sodium chloride. NaCl does not exist as individual molecules; it is a vast, orderly aggregate of $\text{Na}^+$ and $\text{Cl}^-$ ions, an ionic lattice. Compounds formed by such electron transfer are ionic or electrovalent compounds.

Figure

Electron transfer between a sodium atom and a chlorine atom shown as concentric shells, sodium losing its single outer electron to chlorine so that both end up with a stable eight electron outer shell, forming a positive sodium ion and a negative chloride ion

Ionic compounds share properties that trace back to that strong electrostatic attraction. They are hard but brittle — a sharp blow shifts the ion layers so like charges suddenly face each other and repel, shattering the crystal. They have high melting and boiling points (NaCl melts at 1074 K) since pulling the ions apart takes a large amount of energy. Most dissolve readily in water but not in kerosene or petrol. And the ions are locked in place in the solid and cannot carry charge — but once melted or dissolved, they move freely and conduct.

Occurrence and extraction of metals

Elements or compounds occurring naturally in the earth's crust are minerals; one with a high enough metal concentration to extract profitably is an ore. Ores are mixed with unwanted rock, sand and soil, called gangue, removed before extraction. The reactivity series decides the extraction route. Metals at the bottom — gold, silver, platinum, largely copper — are so unreactive they occur native, and some (mercury's ore cinnabar, or copper sulphide ores) need only heating in air. Metals in the middle — zinc, iron, lead, copper — occur mainly as sulphides or carbonates and are first converted to oxides, since oxides reduce more easily: sulphide ores are roasted in excess air; carbonate ores are calcined in limited air.

2\text{ZnS}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{ZnO}(s) + 2\text{SO}_2(g) \qquad \text{(roasting)}

\text{ZnCO}_3(s) \xrightarrow{\Delta} \text{ZnO}(s) + \text{CO}_2(g) \qquad \text{(calcination)}

The oxide is then reduced with carbon, $\text{ZnO}(s) + \text{C}(s) \to \text{Zn}(s) + \text{CO}(g)$ , or, for stubborn oxides such as $\text{Fe}_2\text{O}_3$ , by a more reactive metal such as aluminium in the exothermic thermit reaction, $\text{Fe}_2\text{O}_3(s) + 2\text{Al}(s) \to 2\text{Fe}(l) + \text{Al}_2\text{O}_3(s) + \text{heat}$ , used to weld railway tracks with the molten iron it produces. Metals at the top — potassium, sodium, calcium, magnesium, aluminium — bind oxygen too strongly for carbon to displace, so they need electrolytic reduction of their molten compounds: metal deposits at the cathode, $\text{Na}^+ + e^- \to \text{Na}$ , and non-metal is liberated at the anode, $2\text{Cl}^- \to \text{Cl}_2 + 2e^-$ . Metals from any route are usually impure and undergo electrolytic refining: impure metal as anode, a strip of pure metal as cathode, a solution of the metal's salt as electrolyte — pure metal dissolves off the anode and deposits on the cathode, while insoluble impurities settle as anode mud.

Corrosion and alloys

Corrosion is the slow chemical attack of a metal's exposed surface — silver tarnishes to black silver sulphide, copper greens over to basic copper carbonate, and iron rusts to a flaky brown coating of hydrated iron oxide. Iron rusting needs both air and moisture together; remove either and rust does not form. Prevention blocks that access: painting, oiling, greasing, chrome plating, galvanising (a protective zinc coat), anodising (thickening aluminium's own oxide layer electrolytically), or alloying. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal, made by melting the base metal and dissolving the others in before cooling. It is used because a metal's raw properties are rarely ideal: pure iron is soft, but with about 0.05% carbon it becomes hard steel, and with nickel and chromium it becomes rust-resistant stainless steel. Brass (copper, zinc), bronze (copper, tin) and solder (lead, tin, low melting point, for wiring) are everyday examples; even 22-carat gold jewellery alloys 22 parts gold with 2 parts copper or silver, since pure 24-carat gold is too soft to hold its shape.

Derivation ABuilding the reactivity series from independent experimental evidence

No single reagent can rank every metal, because metals refuse to react with different reagents — copper will not touch water, silver will not touch dilute acid — so any one test leaves large gaps. The reactivity series is pieced together by running several independent tests on the same metals and stitching the results into one consistent ranking, checking that new evidence never contradicts what came before. This is the logical process behind NCERT Activities 3.9 to 3.12, and it is the method, not just the final list, that gets tested.

Step 1. Test reaction with oxygen. Sodium and potassium react so fast that they are stored under kerosene; magnesium burns with a dazzling white flame, less violently; copper tarnishes black with CuO; silver and gold show no reaction. This gives a rough order: Na and K on top, Mg below, Cu lower, Ag and Au essentially unreactive.

Step 2. Test reaction with water, in stages of increasing severity — cold water, hot water, then steam. Sodium and potassium react so violently with cold water that the hydrogen catches fire; calcium reacts gently, floating as bubbles cling to it; magnesium needs hot water; aluminium, iron and zinc need steam; copper, silver and gold never react. This slots calcium, aluminium, iron and zinc into the ranking.

Step 3. Test reaction with dilute hydrochloric acid, timing bubbling rate and temperature rise. Magnesium bubbles fastest, then aluminium, then zinc, then iron, each slower; copper shows no bubbles at all. This sharpens the middle of the series: Mg is greater than Al is greater than Zn is greater than Fe, copper below all of them.

Step 4. Close remaining gaps with displacement: an iron nail in copper(II) sulphate turns reddish-brown as copper deposits, so iron is more reactive than copper, while copper wire in iron(II) sulphate shows nothing. Running this logic across many metal-and-salt pairs resolves every remaining ambiguity, since displacement is unambiguous evidence of relative reactivity even where a metal is inert to oxygen, water and acid. Combining all four steps gives the finished series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, (H), Cu, Hg, Ag, Au, most reactive to least.

Go deeper · Derivation B · 1 minHow sodium and chlorine combine to form the ionic compound sodium chloride

Derivation BHow sodium and chlorine combine to form the ionic compound sodium chloride

A reaction between a metal and a non-metal is, at heart, a story about each atom's outermost shell trying to look like the nearest noble gas, whose shell is completely filled and chemically almost inert. Tracing exactly how sodium and chlorine achieve that shows why the bond forms and why the resulting compound is brittle, high-melting, and never found as a lone molecule.

Step 1. Write each atom's electronic configuration. Sodium (atomic number 11) is 2, 8, 1 — a single lonely electron in the outer M shell. Chlorine (atomic number 17) is 2, 8, 7 — one short of a stable octet.

Step 2. Identify what each atom needs. Sodium reaches a stable, full shell if it discards its one outer electron, letting the already-stable L shell become outermost. Chlorine needs one more electron to complete its octet. The needs are complementary: sodium wants to lose exactly what chlorine wants to gain.

Step 3. Transfer the electron. Sodium loses its outer electron to become $\text{Na}^+$ (configuration 2, 8), a net $+1$ charge with 11 protons but only 10 electrons. Chlorine accepts that electron to become $\text{Cl}^-$ (configuration 2, 8, 8), a net $-1$ charge with 17 protons and 18 electrons. No electron is destroyed; it simply moves.

Step 4. Assemble the compound and read off its properties. The oppositely charged ions attract by strong electrostatic force, and vast numbers pack into a repeating three-dimensional lattice — there is no single "molecule" of NaCl, only the ongoing aggregate. Overcoming that lattice-wide attraction takes a large amount of energy, which is why NaCl melts only at 1074 K, and because the ions are rigidly fixed in the solid, solid NaCl cannot conduct; only once melted or dissolved are they free to carry charge.

Go deeper · Derivation C · 2 minMatching an ore to its extraction route using the activity series

Derivation CMatching an ore to its extraction route using the activity series

Extracting a pure metal from its ore is not one universal recipe; the correct sequence depends on how tightly the metal binds oxygen, sulphur or carbonate in nature, read straight off the reactivity series. Walking through the decision process shows how metallurgists choose between doing almost nothing, reducing with carbon, or resorting to electrolysis.

Step 1. Locate the metal in the series. Near the bottom (copper, silver, mercury, gold), affinity for oxygen or sulphur is so weak the metal may occur free, or its sulphide ore converts straight to metal by heating in air, as with cinnabar (HgS) roasted to HgO and then decomposed to liquid mercury on further heating.

Step 2. In the middle (zinc, iron, lead, copper), strip away the gangue, then convert any sulphide or carbonate ore into an oxide, since oxides reduce more easily. A sulphide ore is roasted in excess air, releasing sulphur dioxide; a carbonate ore is calcined in limited air, releasing carbon dioxide.

Step 3. Reduce the oxide to the free metal. Carbon works well for this middle band — heating zinc oxide with carbon gives zinc and carbon monoxide — but for oxides carbon cannot easily reduce, a more reactive metal such as aluminium is used instead, as in the exothermic thermit reaction, releasing enough heat to produce molten iron on the spot.

Step 4. At the very top of the series (potassium, sodium, calcium, magnesium, aluminium), skip carbon reduction, since these metals bind oxygen more tightly than carbon does. Electrolyse the molten compound instead: the metal ion gains electrons and deposits at the cathode while the non-metal is liberated at the anode, as when molten sodium chloride gives sodium at the cathode and chlorine at the anode.

Step 5. Whatever the route, finish by refining. Metal recovered by roasting, carbon reduction or electrolysis is rarely pure enough for use, so it is commonly purified by electrolytic refining — made the anode in a bath containing a solution of its own salt, with a thin strip of pure metal as the cathode, so that pure metal migrates from anode to cathode while insoluble impurities drop away as anode mud.

Worked examples

Board level - Ranking four unknown metals from a displacement table

Samples of four metals A, B, C and D were each added to solutions of zinc sulphate, iron(II) sulphate, copper(II) sulphate and silver nitrate. The results were: A displaced only iron(II) sulphate; B displaced iron(II) sulphate, copper(II) sulphate and silver nitrate; C displaced none of the four; D displaced only silver nitrate. Rank A, B, C and D in decreasing order of reactivity, and predict what would be seen if metal B were dropped into a solution of copper(II) sulphate.

Step 1. Read each displacement as reactivity evidence: whenever a metal displaces another from its salt solution, the metal doing the displacing is the more reactive of the two.

Step 2. Rank by how many salts each metal displaced. B displaced three of four (iron, copper, silver), placing it highest. A displaced only iron(II) sulphate, placing it below B. D displaced only silver nitrate, the weakest test, placing it below A. C displaced nothing, placing it last, below even silver.

Step 3. Assemble the order and predict the untested case. Decreasing reactivity is B, A, D, C. Since B already displaced copper(II) sulphate, dropping B into copper(II) sulphate again would deposit reddish-brown copper while the blue colour fades. Answer: B is greater than A is greater than D is greater than C, and B in copper(II) sulphate deposits copper as the blue colour fades.

Board level - Electron-dot reasoning for the formation of magnesium oxide

Magnesium (atomic number 12, configuration 2, 8, 2) burns in oxygen (atomic number 8, configuration 2, 6) to form magnesium oxide. Using electron-dot reasoning, show how the compound forms, name the ions present, and give the formula of the compound.

Step 1. Identify each atom's need. Magnesium has two outer electrons and must lose both to reach the stable 2, 8 configuration of neon. Oxygen has six outer electrons and must gain two to complete its octet.

Step 2. Transfer the electrons. Magnesium loses both to become $\text{Mg}^{2+}$ (configuration 2, 8); one oxygen atom gains both to become $\text{O}^{2-}$ (configuration 2, 8). One magnesium atom supplies exactly what one oxygen atom needs, a one-to-one ratio.

Step 3. Write the ionic reaction and formula. $\text{Mg} \to \text{Mg}^{2+} + 2e^-$ and $\text{O} + 2e^- \to \text{O}^{2-}$ , held together by electrostatic attraction. Answer: the ions are $\text{Mg}^{2+}$ and $\text{O}^{2-}$ , and the formula is MgO.

Board HOTS / foundation JEE-NEET - Mass of iron produced in a thermit reaction

In a thermit welding operation, 5.4 g of aluminium powder reacts completely with excess iron(III) oxide according to $\text{Fe}_2\text{O}_3(s) + 2\text{Al}(s) \to 2\text{Fe}(l) + \text{Al}_2\text{O}_3(s) + \text{heat}$ . Given the molar mass of Al is 27 g/mol and of Fe is 56 g/mol, find the mass of molten iron produced.

Step 1. Convert the given mass of aluminium to moles. Moles of Al $= \dfrac{5.4\ \text{g}}{27\ \text{g/mol}} = 0.20\ \text{mol}$ .

Step 2. Use the balanced equation's mole ratio to find moles of iron. The equation shows 2 mol Al produces 2 mol Fe — a one-to-one ratio between Al and Fe — so 0.20 mol Al produces 0.20 mol Fe.

Step 3. Convert moles of iron back to mass. Mass of Fe $= 0.20\ \text{mol} \times 56\ \text{g/mol} = 11.2\ \text{g}$ . Answer: 11.2 g of molten iron is produced. Timing note: about ninety seconds. It is a single mole-ratio bridge — grams to moles, moles to moles via the balanced equation's coefficients, moles back to grams — so the only real trap is reading the mole ratio off the equation correctly rather than assuming it is always 1 is to 1.

Common traps

Treating physical tests as a reliable way to classify every element — mercury, gallium, caesium, iodine, graphite and the soft alkali metals all break the naive rule.
Writing sodium chloride, or any ionic compound, as a discrete molecule — an ionic solid is a continuous lattice, which is why it needs so much energy to melt.
Assuming an ionic compound always conducts electricity — only when molten or dissolved, never as a rigid solid.
Assuming metal oxides are always basic — aluminium oxide and zinc oxide are amphoteric.
Forgetting one reagent does not predict another (aluminium ignores water but reacts vigorously with acid) — the series needs all four tests together.
Believing carbon can reduce every metal oxide — it fails above zinc, where metals bind oxygen too strongly and need electrolysis instead.
Assuming a reactive metal is unsafe to use — aluminium sits well above iron yet makes excellent cooking vessels, protected by its own instant oxide layer.
Confusing roasting (sulphide, excess air, $\text{SO}_2$ ) with calcination (carbonate, limited air, $\text{CO}_2$ ).

Common trapIonic compounds do not exist as individual molecules

The single most common misconception in this chapter is picturing sodium chloride, or any ionic compound, as a self-contained molecule of one cation stuck to one anion, the way $\text{H}_2\text{O}$ or $\text{CO}_2$ are genuine molecules. In reality, an ionic solid is one continuous three-dimensional lattice in which every $\text{Na}^+$ ion is surrounded by several $\text{Cl}^-$ ions and vice versa, held together by electrostatic attraction extending through the entire crystal, not confined to one pair of ions. This is why ionic compounds have such high melting points compared with molecular compounds of similar size — melting means overcoming the cumulative attraction of the whole lattice, not breaking one bond per molecule — and why the "formula unit" NaCl is a ratio of ions in the lattice, not an isolated molecule.

You should be able to…

List the physical properties distinguishing metals from non-metals, with an exception to each.
Write balanced equations for a metal's reaction with oxygen, water and dilute acid.
Explain what an amphoteric oxide is and name two examples.
Use the reactivity series to predict whether a displacement reaction will occur.
Explain, using electron transfer and shell configurations, how an ionic compound such as NaCl or MgO forms.
List the properties of ionic compounds and explain each in terms of the ionic lattice.
Match a metal's reactivity-series position to its extraction route: native, carbon reduction after roasting/calcination, or electrolysis.
Explain the conditions needed for iron to rust, and three practical ways to prevent corrosion.

Ask the tutor: if you are unsure why aluminium is both highly reactive and perfectly safe for cooking, or want to practise ranking unfamiliar metals from a displacement table, ask your tutor to work through one with you.

Aluminium is high up in the reactivity series, so why is it safe to cook in an aluminium pan?

Aluminium reacts with oxygen the moment it meets air, but that reaction protects it: the metal instantly forms a thin, tough oxide layer that sticks tightly and stops oxygen or moisture reaching the aluminium underneath. So the pan is not unreactive — its surface reacted once, permanently, and now shields the rest of the metal. Anodising uses the same idea deliberately, thickening that natural oxide layer electrolytically for even more corrosion resistance.

Why is 22-carat gold used for jewellery instead of pure 24-carat gold?

Pure gold (24 carat) is soft enough that fine jewellery made from it would bend and scratch too easily for daily wear. Alloying it with a little copper or silver makes it noticeably harder while barely changing its colour. In 22-carat gold, 22 parts out of 24 are pure gold and the remaining 2 parts are copper or silver — which is why 22-carat jewellery is both more durable and slightly cheaper than the same weight of pure gold.

What exactly is the difference between roasting and calcination?

Both convert an ore into the metal oxide by strong heating, but from different ores and different air conditions. Roasting treats sulphide ores in an excess of air, so sulphur burns away as sulphur dioxide gas, leaving the oxide behind. Calcination treats carbonate ores with limited air, driving off carbon dioxide gas. The gas released, sulphur dioxide versus carbon dioxide, is the quickest way to tell them apart in an exam answer.

Why can carbon reduce zinc oxide but not sodium oxide or aluminium oxide?

Carbon reduction works only when carbon's own attraction for oxygen is stronger than the metal's, so carbon can "steal" the oxygen and leave the free metal behind. Zinc, iron and lead hold oxygen less strongly than carbon does, so carbon reduction works for them. But sodium, magnesium, calcium and aluminium hold oxygen far more strongly than carbon — it simply cannot pull the oxygen away, however hard it is heated. Those metals need electrolytic reduction instead, where electrical energy forces the metal ion to accept electrons and become the free metal.