1. Introduction – Metals and Non-metals
Around you, every material can be placed in one of two big families in chemistry: metals and non-metals. Your steel lunch box, copper wire, aluminium foil and iron tools are metals. The oxygen you breathe, the carbon in pencil lead, the phosphorus on a matchbox and the sulphur used in fungicides are non-metals.
A metal is an element that generally has lustre, is malleable and ductile, is a good conductor of heat and electricity, is sonorous, and usually forms basic oxides. Examples: iron (Fe), copper (Cu), gold (Au), sodium (Na).
A non-metal is an element that generally lacks metallic lustre, is brittle or soft, is a poor conductor of heat and electricity, is non-sonorous, and usually forms acidic oxides. Examples: oxygen (O), carbon (C), sulphur (S), chlorine (Cl).
Occurrence of metals and non-metals
Most metals are not found free in nature; they occur combined as oxides, sulphides, carbonates and other ores in the earth's crust. An important exception is the group of less reactive metals, such as gold and silver, that can occur in a free (native) state. Non-metals occur in a wide variety of forms: oxygen and nitrogen as gases of the air, carbon as coal and graphite, and sulphur as underground deposits.
Why metals and non-metals matter
- Metals give us tools, vehicles, electrical wires, jewellery, utensils and construction beams.
- Non-metals are essential for life (oxygen, nitrogen) and for industry (sulphuric acid from sulphur, fertilisers from nitrogen).
- Almost every modern device uses both — the copper wire inside a phone carries the electricity produced from carbon-based fuels or batteries.
Fewer than 25 elements are non-metals; the vast majority of the 118 elements are metals. But oxygen, nitrogen and carbon — all non-metals — make up most of the living matter and the air, so non-metals are just as important in everyday chemistry.
2. Physical Properties of Metals
Metals can usually be recognised by a set of physical properties. These properties are general — a few metals have exceptions, so think of them as typical, not universal, rules.
Lustre
Metals have a shiny appearance called metallic lustre. Body parts such as gold and silver polish to a bright shine. Their surfaces reflect light strongly because the delocalised free electrons vibrate and reflect light waves.
Hardness
Most metals are hard and strong, which is why they are used to make tools, machinery and structures. The strongest metals resist scratching and bending.
Malleability
The property of being beaten or pressed into thin sheets without breaking (without cracking). Example: aluminium foil, gold leaves. Malleability relies on the ability of metal atoms to slide over each other under pressure.
Ductility
The property of being drawn (stretched) into thin wires. Example: copper and aluminium are drawn into electrical wires. A ductile metal can be pulled into a long, thin thread.
Conductivity of heat and electricity
Metals are good conductors of heat, so they warm up quickly and cool quickly — that is why metal pans are used for cooking. They are also good conductors of electricity, which is why wires are made of copper and aluminium. Silver is the best conductor and is used in special electronics despite its cost.
Sonority
The property of producing a ringing sound when struck. Metal bells and gongs make a clear, musical sound. Non-metals are non-sonorous: sulphur or carbon makes a dull thud when struck.
Melting and boiling points
Metals generally have high melting and boiling points. Iron melts at about 1539°C and copper at about 1085°C. This is why metal objects do not lose their shape in ordinary heat.
Physical state
At room temperature, almost all metals are solids. The most important exception is mercury (Hg), which is a liquid at room temperature — the only metal that is liquid.
| Property | Metals | Example / exception |
|---|---|---|
| Lustre | Shiny / lustrous | Gold, silver |
| Hardness | Generally hard | Exceptions: sodium and potassium are soft enough to be cut with a knife |
| Malleability | Can be hammered into sheets | Aluminium foil, gold leaf |
| Ductility | Can be drawn into wires | Copper wire, aluminium wire |
| Electrical conductivity | Good (solid and molten) | Copper, silver, aluminium |
| Thermal conductivity | Good | Metal pans for cooking |
| Sonority | Sonorous (ringing sound) | Bells |
| Melting / boiling point | Generally high | Exception: gallium melts in the palm (about 30°C) |
| Physical state (room temperature) | Mostly solid | Exception: mercury is a liquid |
Do not say "all metals are hard solids". Sodium and potassium are soft, mercury is a liquid, and gallium melts easily. Mention these exceptions when a question asks for exceptions.
3. Physical Properties of Non-Metals
Non-metals are different in physical character from metals in almost every respect: they are mostly dull, brittle, poor conductors and non-sonorous. As with metals, exceptions exist (for example, graphite conducts electricity and has a slippery, shiny graphite form), so keep the general picture in mind.
Appearance
Non-metals generally have a dull appearance. Carbon (charcoal), sulphur and phosphorus are not shiny. (Diamond, a form of carbon, is the exception that polishes to brilliance.)
Brittleness and hardness
Non-metals are usually brittle in the solid state — they break or crumble when hammered; they cannot be drawn into wires (non-ductile) nor beaten into sheets (non-malleable). Sulphur and phosphorus powders are typical. Exception: diamond is the hardest known natural material but it is still brittle; it does not flow or stretch.
Conductivity
Non-metals are generally poor conductors of heat and electricity — they are insulators. Air, glass (inert), sulphur and plastic allow no current. Exception: graphite (a form of carbon) conducts electricity, which is why it is used in electrodes, pencils and lubricants (its layers slide over each other).
Sonority
Non-metals are non-sonorous: they produce no ringing sound when struck.
States and melting points
Non-metals exist in all three states at room temperature: gases (oxygen, hydrogen, nitrogen, chlorine), liquid (bromine) and solids (carbon, sulphur, phosphorus). Non-metals generally have low melting and boiling points (diamond is a striking exception, with a very high melting point due to its giant structure).
| Non-metal | State at room temperature | Typical use |
|---|---|---|
| Oxygen (O) | Gas | Respiration, combustion |
| Nitrogen (N) | Gas | Fertilisers, inert filling |
| Hydrogen (H) | Gas | Fuels, ammonia |
| Chlorine (Cl) | Gas | Water purification, bleaching |
| Bromine (Br) | Liquid | Laboratory reagent |
| Carbon (C) | Solid | Fuel (coal), electrodes (graphite), diamonds |
| Sulphur (S) | Solid | Sulphuric acid, matches, fungicides |
| Phosphorus (P) | Solid | Matchboxes |
Metals vs Non-metals – comparison table
| Property | Metal | Non-metal |
|---|---|---|
| Distinct appearance | Lustrous / shiny | Dull |
| Malleability | Malleable (sheet on hammer) | Non-malleable (brittle) |
| Ductility | Ductile (wires) | Non-ductile |
| Hardness | Generally hard (Na, K are soft) | Generally brittle or soft (diamond is very hard) |
| Heat / electricity | Good conductor | Poor (except graphite for electricity) |
| Sonority | Sonorous (ring) | Non-sonorous (thud) |
| State at RT | Solid (Hg liquid) | Gas, liquid (Br), solid |
| Melting / boiling point | Generally high | Generally low |
| Nature of oxide | Basic (amphoteric Al, Zn) | Acidic |
| Ions formed | Positive (cation) — electrons lost | Negative (anion) — electrons gained |
| Examples | Fe, Cu, Au, Na, Al, Hg | O, N, C, S, P, Cl, H |
Metals are shiny, bendable, stretchable, sonorous, conductive and form basic oxides. Non-metals are the opposite: dull, brittle, non-sonorous, insulating and form acidic oxides. Learn the three famous exceptions as a package: mercury (liquid metal), iodine (shiny non-metal), graphite (conducting non-metal).
4. Chemical Properties of Metals
Metals react by losing electrons. Their reactivity depends on how readily they lose electrons — which is exactly the idea captured by the reactivity series (Section 6).
A. Metals with oxygen
Most metals combine with oxygen of the air to form metal oxides. Metal oxides are basic oxides — they react with acids to give salt and water, and dissolve in water to give bases.
| Reaction | Balanced equation | Observation |
|---|---|---|
| Sodium + oxygen | \[ 4Na(s) + O_2(g) \xrightarrow{\Delta} 2Na_2O(s) \] | Yellow-white sodium peroxide / oxide layer forms on sodium |
| Magnesium + oxygen | \[ 2Mg(s) + O_2(g) \xrightarrow{\Delta} 2MgO(s) \] | Dazzling white flame; white magnesium oxide ash |
| Calcium + oxygen | \[ 2Ca(s) + O_2(g) \xrightarrow{\Delta} 2CaO(s) \] | Quicklime (CaO) forms |
| Aluminium + oxygen | \[ 4Al(s) + 3O_2(g) \xrightarrow{\Delta} 2Al_2O_3(s) \] | Fine aluminium powder burns brightly |
| Copper + oxygen | \[ 2Cu(s) + O_2(g) \xrightarrow{\Delta} 2CuO(s) \] | Black copper(II) oxide forms on the surface |
Basic oxides
Basic oxides are metal oxides that react with acids to form salt and water, and turn red litmus blue in aqueous solution. Na₂O, MgO and CaO dissolve in water to give alkalis:
\[ Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq) \]
\[ CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) \]
Amphoteric oxides
An oxide that behaves as a basic oxide (reacts with acids) and also as an acidic oxide (reacts with bases). Aluminium oxide and zinc oxide are amphoteric.
\[ Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l) \]
\[ Al_2O_3(s) + 2NaOH(aq) \rightarrow 2NaAlO_2(aq) + H_2O(l) \]
B. Metals with water
Metals differ a lot in how they react with water. The more reactive metals react quickly with cold water; less reactive ones need steam; and least reactive ones do not react at all. This difference is controlled by the reactivity series.
| Metal | Reaction with water | Balanced equation |
|---|---|---|
| Potassium (K) | Violent reaction with cold water; hydrogen burns | \[ 2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g) \] |
| Sodium (Na) | Reacts vigorously with cold water, floats and moves; hydrogen catches fire | \[ 2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g) \] |
| Calcium (Ca) | Reacts less violently with cold water; bubbles of hydrogen | \[ Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g) \] |
| Magnesium (Mg) | Barely reacts with cold water; reacts with hot water and vigorously with steam | \[ Mg(s) + H_2O(g) \xrightarrow{\Delta} MgO(s) + H_2(g) \] |
| Aluminium (Al) | Does not react with cold or hot water; reacts with steam | \[ 2Al(s) + 3H_2O(g) \xrightarrow{\Delta} Al_2O_3(s) + 3H_2(g) \] |
| Zinc (Zn) | Does not react with cold water; reacts with steam | \[ Zn(s) + H_2O(g) \xrightarrow{\Delta} ZnO(s) + H_2(g) \] |
| Iron (Fe) | Does not react with cold water; reacts with steam (reversible) | \[ 3Fe(s) + 4H_2O(g) \xrightarrow{\Delta} Fe_3O_4(s) + 4H_2(g) \] |
| Copper (Cu) | Does not react with water or steam | No reaction |
Metal + water does not always give metal hydroxide + hydrogen. For less reactive metals (Mg with steam, Fe with steam), the product is the metal oxide + hydrogen. And copper, silver, gold and mercury do not react with water at all.
C. Metals with dilute acids
More reactive metals displace hydrogen from dilute acids, producing a salt and hydrogen gas.
General pattern: Metal + Dilute acid → Salt + Hydrogen gas
| Reaction | Balanced equation | Observation |
|---|---|---|
| Zinc + dilute sulphuric acid | \[ Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g) \] | Bubbles of hydrogen; zinc dissolves |
| Iron + dilute hydrochloric acid | \[ Fe(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2(g) \] | Hydrogen evolves; light green ferrous chloride forms |
| Magnesium + dilute hydrochloric acid | \[ Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g) \] | Fast and effervescent; Mg dissolves quickly |
| Aluminium + dilute hydrochloric acid | \[ 2Al(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2(g) \] | Hydrogen evolves (slow at first due to oxide layer) |
| Copper + dilute HCl / H₂SO₄ | No reaction | Copper is below hydrogen in the series |
Testing for hydrogen
Bring a burning splinter near the mouth of the tube. Hydrogen burns with a pop sound (it is combustible), confirming hydrogen gas.
D. Metals with salt solutions (displacement)
A more reactive metal can displace a less reactive metal from its salt solution. This is a displacement reaction (Section 7 gives full treatment).
Iron (more reactive) displaces copper from a solution of copper sulphate. Blue copper sulphate solution fades, red-brown copper deposits on the iron nail.
\[ Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s) \]
E. Metals with bases
Most metals do not react with bases. However, some amphoteric metals such as zinc and aluminium react with strong alkalis (sodium hydroxide) to give hydrogen gas and a metal salt.
| Reaction | Balanced equation | Observation |
|---|---|---|
| Zinc + sodium hydroxide solution | \[ Zn(s) + 2NaOH(aq) \rightarrow Na_2ZnO_2(aq) + H_2(g) \] | Hydrogen gas; sodium zincate solution |
| Aluminium + sodium hydroxide solution | \[ 2Al(s) + 2NaOH(aq) + 2H_2O(l) \rightarrow 2NaAlO_2(aq) + 3H_2(g) \] | Hydrogen gas; sodium aluminate solution |
5. Chemical Properties of Non-Metals
Non-metals react by gaining electrons (forming negative ions/compounds) or by forming covalent bonds. Their reaction patterns differ from metals in important ways.
A. Reaction with oxygen – acidic oxides
Non-metals combine with oxygen to form non-metal oxides, which are acidic oxides. These dissolve in water to form acids and neutralise bases.
| Reaction | Balanced equation | Observation / property |
|---|---|---|
| Carbon (charcoal) + oxygen | \[ C(s) + O_2(g) \xrightarrow{\Delta} CO_2(g) \] | CO₂ is acidic; turns lime water milky |
| Sulphur + oxygen | \[ S(s) + O_2(g) \xrightarrow{\Delta} SO_2(g) \] | Choking gas; SO₂ forms sulphurous acid in water |
| Phosphorus + oxygen | \[ 4P(s) + 5O_2(g) \xrightarrow{\Delta} 2P_2O_5(s) \] | Dense white fumes of phosphorus pentoxide |
Acidic oxides react with bases — that is why CO₂ turns lime water milky:
\[ Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l) \]
B. Reaction with hydrogen
Some non-metals react with hydrogen to form stable covalent hydrides (compounds).
| Reaction | Balanced equation | Property |
|---|---|---|
| Chlorine + hydrogen | \[ H_2(g) + Cl_2(g) \xrightarrow{sunlight} 2HCl(g) \] | Colourless gas; forms hydrochloric acid in water |
| Nitrogen + hydrogen | \[ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \] | Ammonia — used for fertilisers (Haber process) |
| Sulphur + hydrogen | \[ H_2(g) + S(s) \xrightarrow{\Delta} H_2S(g) \] | Hydrogen sulphide; rotten-egg smell |
C. Reaction with metals
Metals lose electrons to non-metals, forming ionic compounds. For example:
\[ 2Na(s) + Cl_2(g) \xrightarrow{\Delta} 2NaCl(s) \]
\[ 2Fe(s) + 3Cl_2(g) \xrightarrow{\Delta} 2FeCl_3(s) \]
Sodium loses one electron to chlorine to give Na⁺ and Cl⁻, which attract to form the ionic compound sodium chloride.
D. Non-metals do not displace hydrogen from dilute acids
Unlike active metals, a non-metal cannot displace hydrogen from a dilute acid (there is no metal ion to replace). Therefore non-metals as such do not react with dilute HCl or dilute H₂SO₄ to give hydrogen gas.
Two quick summary rules: (1) metal oxides are basic, non-metal oxides are acidic; (2) metals form cations (Na⁺, Ca²⁺, Al³⁺), non-metals form anions (Cl⁻, O²⁻, S²⁻).
6. Reactivity Series of Metals
Not all metals react equally. Sodium fizzes violently in water while gold does not react at all. The reactivity series is the list of metals arranged from most reactive to least reactive based on how readily they lose electrons.
The arrangement of metals in the decreasing order of their reactivity (ease of electron loss and reaction with water, acids, etc.) is called the reactivity series. Hydrogen is placed in the series for comparison even though it is a non-metal.
The standard series (decreasing reactivity)
| Position | Metal | Reaction with water | Reaction with dilute acid (HCl/H₂SO₄) | Extraction clue |
|---|---|---|---|---|
| Most reactive | K | Violent with cold water | React vigorously with dilute acids — liberate H₂ | Extracted by electrolysis (high/near top) |
| Na | Vigorous with cold water | |||
| Ca | Reacts with cold water | |||
| Mg | Slow with cold, fast with steam | |||
| Reactive | Al | React with steam | ||
| Zn | Steam only | |||
| Fe | Steam only | |||
| Pb | Steam only (slow) | |||
| — | H (for comparison) | — | — | — |
| Least reactive | Cu | No reaction | Do not react with dilute acids (below hydrogen) | Found native or extracted by roasting/calcination |
| Hg | No reaction | |||
| Ag | No reaction | |||
| Au | No reaction |
\[ K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au \]
What the series tells us
- Most reactive (K, Na, Ca): react violently with cold water, release hydrogen, and are stored suitably (sodium in kerosene).
- Medium reactivity (Mg, Al, Zn, Fe, Pb): react with steam and dilute acids, form stable oxide ores, and are extracted by reduction of their oxides with carbon.
- Least reactive (Cu, Hg, Ag, Au): do not react with dilute acids; their compounds are easily decomposed by heat alone, and Ag/Au occur free in nature.
- Displacement reactions and extraction processes are both predicted from the position in the series.
Students often place Iron above Zinc. The correct order is Al > Zn > Fe. Also remember hydrogen sits between lead and copper — metals above H displace H₂ from acids; metals below H do not.
7. Displacement Reactions
A reaction in which a more reactive element displaces (takes the place of) a less reactive element from its compound. In a metal + salt-solution context: More reactive metal + salt of less reactive metal → its own salt + less reactive metal (as free metal).
Because the more reactive metal loses electrons more easily, it removes the less reactive metal from its compound. This is exactly what the position in the reactivity series predicts: a metal higher up displaces a metal lower down from its salt solution.
Examples with explanations
| No. | Reaction | Equation | Observation |
|---|---|---|---|
| 1 | Iron + copper sulphate | \[ Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s) \] | Blue solution fades; brown copper deposits |
| 2 | Zinc + copper sulphate | \[ Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s) \] | Blue solution fades; Cu deposits |
| 3 | Copper + silver nitrate | \[ Cu(s) + 2AgNO_3(aq) \rightarrow Cu(NO_3)_2(aq) + 2Ag(s) \] | Silver crystals deposit on the wire; solution turns blue |
| 4 | Zinc + ferrous sulphate | \[ Zn(s) + FeSO_4(aq) \rightarrow ZnSO_4(aq) + Fe(s) \] | Light green solution turns colourless as Fe deposits |
| 5 | Aluminium + copper(II) chloride | \[ 2Al(s) + 3CuCl_2(aq) \rightarrow 2AlCl_3(aq) + 3Cu(s) \] | Blue-green solution fades; copper deposits |
| 6 | Magnesium + silver nitrate | \[ Mg(s) + 2AgNO_3(aq) \rightarrow Mg(NO_3)_2(aq) + 2Ag(s) \] | Grey silver deposits; Mg dissolves |
When no displacement happens
If the added metal is less reactive than the metal in the salt, no reaction occurs. For example, copper dipped in a solution of zinc sulphate or ferrous sulphate shows no change, because Cu is below Zn and Fe in the series.
\[ Cu(s) + ZnSO_4(aq) \rightarrow \text{no reaction} \]
Reactivity depends on how easily an atom loses electrons to form a cation. K loses electrons far more readily than Au. In a salt solution, the more reactive metal donates its electrons to the less reactive metal's ions, which settle out as free metal — hence displacement.
In exams, memorise the series, then answer "which displaces which" by comparing positions. A metal can never displace a metal above it in the series.
8. Ionic Compounds
When a metal reacts with a non-metal, electrons are transferred from the metal atoms to the non-metal atoms. The metal becomes a positively charged cation; the non-metal becomes a negatively charged anion. The full transfer of electrons forms an ionic (or electrovalent) compound.
Formation of an ionic compound – the NaCl model
- Sodium (Na, atomic number 11) has 11 electrons: 2, 8, 1 (one electron in the outer shell).
- Chlorine (Cl, atomic number 17) has 17 electrons: 2, 8, 7 (seven electrons in the outer shell).
- Sodium loses its one outer electron to become a sodium cation:
\[ Na \rightarrow Na^+ + e^- \]
- Chlorine gains that electron to become a chloride anion:
\[ Cl + e^- \rightarrow Cl^- \]
- The oppositely charged Na⁺ and Cl⁻ attract each other by strong electrostatic forces; the sodium and chloride ions pack into a giant lattice. Electrons transferred from metal to non-metal form the ionic bond.
Cations, anions and electrovalency
- A cation is a positively charged ion formed when a metal loses electrons (Na⁺, Ca²⁺, Al³⁺).
- An anion is a negatively charged ion formed when a non-metal gains electrons (Cl⁻, O²⁻, S²⁻).
- Electrovalency is the number of electrons lost or gained by an atom in the formation of an ionic compound. Sodium has an electrovalency of 1 (loses 1 electron), calcium has 2, aluminum has 3; chlorine has 1 (gains 1).
Ca loses 2 electrons (Ca²⁺) and chlorine gains 1 each (Cl⁻). To balance charge, two Cl⁻ ions combine with one Ca²⁺: calcium chloride has formula CaCl₂. Similarly Al³⁺ with O²⁻ gives Al₂O₃.
Properties of ionic compounds
| Property | Explanation |
|---|---|
| Solid state | Ionic compounds are crystalline solids at room temperature; ions are held tightly in a regular lattice. |
| Hard and brittle | Strong electrostatic forces make them hard; on hammering the lattice slips and repels, so they break. |
| High melting and boiling points | A large amount of energy is needed to overcome the strong ionic bonds. |
| Soluble in water / insoluble in organic solvents | Water separates the ions (their charges attract polar water molecules); organic solvents do not separate the ions well. |
| Conduction of electricity: solid | Does not conduct — ions are fixed in the lattice and cannot move. |
| Conduction: molten or aqueous | Conducts electricity — ions become free to move and carry charge. |
In solid sodium chloride the ions occupy fixed positions in the crystal lattice and cannot move, so no charge flows. When NaCl melts (about 801°C) or dissolves in water, the lattice breaks and Na⁺ and Cl⁻ ions become mobile; these moving charged ions carry the electric current.
Students write "ionic compounds conduct electricity in the solid state too." They do not — free-moving ions are needed. Conduction happens only in the molten or aqueous state (or if the solid contains mobile ions, which is not the case for NaCl-type lattices).
9. Occurrence of Metals
Metals are extracted from minerals and ores found in the earth's crust. Three terms matter in this section: mineral, ore and gangue.
A naturally occurring material in the earth's crust in which a metal (or its compound) is present. Example: aluminium is present in clay as well as bauxite, so both are minerals of aluminium.
A mineral from which a metal can be extracted economically and profitably. All ores are minerals, but not all minerals are ores. Example: bauxite (Al₂O₃·2H₂O) is the ore of aluminium, while clay is not an ore because extraction is not economical.
The unwanted earthy impurities (sand, rocky material, lime) that accompany the ore in a naturally mined mineral.
Mineral vs Ore
| Basis | Mineral | Ore |
|---|---|---|
| Definition | Natural compound/source of a metal | Mineral from which metal is extracted profitably |
| Economy | May or may not be economical | Always economical to extract |
| Relation | Parent category | Subset of minerals |
| Example | Clay (aluminium mineral) | Bauxite (aluminium ore) |
Reactivity and occurrence
- Least reactive metals (Au, Ag, Cu): often found free/native in nature, or in ores that are easily decomposed.
- Medium reactive metals (Zn, Fe, Pb): generally found as oxides, carbonates or sulphides; the metal has combined with non-metals.
- Very reactive metals (Na, Ca, Al, Mg, K): occur only combined, usually as chlorides, oxides and other compounds; they are extracted by electrolysis.
Common ores (Class 10 examples)
| Metal | Ore | Formula |
|---|---|---|
| Sodium | Common salt (rock salt) | NaCl |
| Aluminium | Bauxite | Al₂O₃·2H₂O |
| Zinc | Zinc blende / zinc carbonate | ZnS / ZnCO₃ |
| Iron | Haematite | Fe₂O₃ |
| Copper | Copper pyrites | CuFeS₂ |
| Mercury | Cinnabar | HgS |
10. Extraction of Metals
The complete set of processes used to extract a metal from its ore, including concentration, extraction (reduction) and refining.
The method chosen depends almost entirely on the reactivity of the metal. Very reactive metals need strong methods (electrolysis); less reactive ones need only heating. In every case the ore is first concentrated (gangue is removed) before extraction.
A. Metals of low reactivity – heating (roasting) or occurring native
Low reactivity metals (Hg, Cu, Ag, Au) have oxides that are easily reduced. Often the ore can be decomposed just by strong heating in air.
Cinnabar is roasted in air to give the metal directly (roasting of sulphide); any mercury oxide formed decomposes on further strong heating:
\[ 2HgS(s) + 3O_2(g) \xrightarrow{\Delta} 2HgO(s) + 2SO_2(g) \]
\[ 2HgO(s) \xrightarrow{\Delta} 2Hg(l) + O_2(g) \]
(Some board-level references write a single combined step for the roasting of cinnabar: \[ HgS(s) + O_2(g) \xrightarrow{\Delta} Hg(l) + SO_2(g) \] — the two-step form above shows the chemistry clearly.)
Gold and silver occur free (native) in nature, so they need no chemical reduction — they are separated by physical methods and then purified.
B. Metals of medium reactivity – roasting / calcination, then reduction by carbon (or CO)
Metals like zinc, iron and lead react too strongly to be extracted by simple heating, but not so strongly as to need electrolysis. Their ores are first converted to the metal oxide, which is then reduced with carbon (or carbon monoxide).
Roasting
Heating an ore (especially a sulphide) in a plentiful supply of air (excess oxygen) to convert it into the metal oxide (and drive off SO₂).
\[ 2ZnS(s) + 3O_2(g) \xrightarrow{\text{roasting}} 2ZnO(s) + 2SO_2(g) \]
Calcination
Heating an ore (usually a carbonate or hydroxide) in a limited supply of air (or in its absence) to drive off water, carbon dioxide or volatile impurities and convert it into the oxide.
\[ ZnCO_3(s) \xrightarrow{\text{calcination}} ZnO(s) + CO_2(g) \]
Reduction by carbon (smelting) and by carbon monoxide
\[ ZnO(s) + C(s) \xrightarrow{\Delta} Zn(s) + CO(g) \]
\[ Fe_2O_3(s) + 3CO(g) \xrightarrow{\Delta} 2Fe(s) + 3CO_2(g) \]
Roasting uses a sulphide ore + oxygen. Calcination uses a carbonate/hydroxide ore with no/controlled air. Then the oxide + carbon (or CO) gives the metal — this is reduction.
C. Metals of high reactivity – electrolytic reduction
Sodium, calcium, magnesium and aluminium are so reactive that their oxides cannot be reduced by carbon. Instead the molten (fused) ore is broken down by passing electricity — electrolysis. The metal forms at the cathode.
\[ 2NaCl(l) \xrightarrow{electricity} 2Na(s) + Cl_2(g) \]
\[ 2Al_2O_3(l) \xrightarrow{electricity} 4Al(l) + 3O_2(g) \]
Do not say "iron and zinc are extracted by electrolysis". Only the most reactive metals (above aluminium; also Al itself) are extracted by electrolysis of molten compounds. Medium-reactive metals are obtained by reducing their oxides with carbon/CO.
11. Refining of Metals
Even after the metal is obtained from its ore, it is usually impure. The process of obtaining a pure metal from its impure form is called refining. The most common method is electrolytic refining, and copper is the classic Class 10 example.
A process in which impure metal is dissolved (from the anode) and pure metal is deposited at the cathode while an electric current flows through an electrolyte containing the metal's salt.
The setup for copper
| Part | Material used | Role |
|---|---|---|
| Anode (+) terminal | Impure copper block | Impure copper dissolves and goes into the solution as Cu²⁺ ions |
| Cathode (–) terminal | Thin sheet of pure copper | Pure copper deposits on it as Cu²⁺ ions gain electrons |
| Electrolyte | Acidified copper sulphate solution (CuSO₄ + dilute H₂SO₄) | Carries Cu²⁺ ions from anode to cathode |
What happens
- At the anode: impure copper dissolves — Cu → Cu²⁺ + 2e⁻. Insoluble impurities (silver, gold, platinum) settle below the anode as anode mud.
- The Cu²⁺ ions travel through the electrolyte to the cathode.
- At the cathode: Cu²⁺ + 2e⁻ → Cu, depositing a layer of pure copper.
- Anode mud is collected and sold — the "impurities" include valuable silver and gold.
Remember the three W's: Work-piece (impure copper) is the anode; pure copper is plated on the cathode; the electrolyte is copper sulphate with dilute acid. And anode mud contains silver and gold — the by-products sold to pay for the process.
12. Corrosion
The gradual wearing away (deterioration) of a metal due to reaction with air, moisture (water), acids, salts or other substances in its surroundings.
Three common examples:
- Rusting of iron — a brown, flaky iron oxide (rust) forms.
- Copper — develops a green coating of basic copper carbonate over time.
- Silver — turns black due to a coating of silver sulphide (Ag₂S).
Rusting of iron – conditions
Rusting of iron requires both air (oxygen) and moisture (water). If either is absent, rusting does not occur. The rust is hydrated iron(III) oxide.
\[ 4Fe(s) + 3O_2(g) + 2xH_2O(l) \rightarrow 2Fe_2O_3 \cdot xH_2O(s) \]
This is the one equation for rusting worth learning. Experiments show that iron in dry air, or completely under oil (no water), does not rust, but iron in normal damp air rusts.
Prevention of corrosion
| Method | How it works | Everyday example |
|---|---|---|
| Painting | A paint film separates the metal from air and moisture | Painted iron gates, window grills, bridges |
| Oiling / greasing | Oil or grease film blocks air and moisture | Cycle chains, machine parts |
| Galvanisation | Coating iron with a layer of zinc metal (zinc protects by being sacrificed and by forming a barrier) | Galvanised iron (GI) sheets, water pipes, buckets |
| Tinning / chrome plating | Coating with tin or chromium | Food cans (tin), tap fittings (chrome) |
| Alloying | Adding other metals to change properties and resist corrosion | Stainless steel kitchenware |
Zinc is more reactive than iron. If the zinc coating is scratched, the zinc itself corrodes (oxidises) first, while the iron is kept safe. This is an example of sacrificial protection; painting would not give this benefit.
Rusting needs both oxygen and water — students forget "moisture" and write rust forms in dry air. Also, corrosion of copper gives a green basic carbonate coating, not simply black rust.
13. Alloys
A homogeneous mixture (solid solution) of a metal with one or more other metals (or with a non-metal, such as carbon in steel), made to improve the properties of the pure metal.
Why alloys are made
- Increase hardness — pure gold or copper is too soft for many uses.
- Prevent corrosion — stainless steel resists rusting.
- Lower the melting point — solder (Pb+Sn) melts easily for joining.
- Improve strength and durability for bridges, tools and machinery.
Important alloys and their constituents/uses
| Alloy | Major constituents | Important use |
|---|---|---|
| Brass | Copper + zinc | Musical instruments, decorative items, taps |
| Bronze | Copper + tin | Statues, medals, bells |
| Solder | Lead + tin | Joining electrical wires and metal parts (low melting point) |
| Steel | Iron + carbon | Tools, construction, rails |
| Stainless steel | Iron + chromium + nickel (+ carbon) | Cutlery, kitchenware, surgical instruments |
| Amalgam | Mercury + another metal | Dental fillings |
Solder is the classic "lower melting point" alloy, and stainless steel is the classic "corrosion-resistant" alloy. Revision tables like the one above often settle one-mark questions.
14. Important Chemical Equations
| No. | Reaction | Balanced equation | Type / Concept |
|---|---|---|---|
| 1 | Sodium + oxygen | \[ 4Na + O_2 \xrightarrow{\Delta} 2Na_2O \] | Metal + O₂ → basic oxide |
| 2 | Magnesium + oxygen | \[ 2Mg + O_2 \xrightarrow{\Delta} 2MgO \] | Bright white flame; basic oxide |
| 3 | Aluminium + oxygen | \[ 4Al + 3O_2 \xrightarrow{\Delta} 2Al_2O_3 \] | Amphoteric oxide |
| 4 | Sodium + water | \[ 2Na + 2H_2O \rightarrow 2NaOH + H_2 \] | Vigorous; H₂ given off |
| 5 | Calcium + water | \[ Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2 \] | Less violent than Na |
| 6 | Iron + steam | \[ 3Fe + 4H_2O(g) \xrightarrow{\Delta} Fe_3O_4 + 4H_2 \] | Metal oxide + H₂ (steam) |
| 7 | Zinc + dilute H₂SO₄ | \[ Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2 \] | Metal + acid → salt + H₂ |
| 8 | Iron + dilute HCl | \[ Fe + 2HCl \rightarrow FeCl_2 + H_2 \] | Pop test for H₂ |
| 9 | Iron + copper sulphate | \[ Fe + CuSO_4 \rightarrow FeSO_4 + Cu \] | Displacement (Fe above Cu) |
| 10 | Copper + silver nitrate | \[ Cu + 2AgNO_3 \rightarrow Cu(NO_3)_2 + 2Ag \] | Displacement (Cu above Ag) |
| 11 | Zinc + NaOH | \[ Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2 \] | Amphoteric metal + base |
| 12 | Carbon + oxygen | \[ C + O_2 \xrightarrow{\Delta} CO_2 \] | Non-metal + O₂ → acidic oxide |
| 13 | Hydrogen + chlorine | \[ H_2 + Cl_2 \xrightarrow{sunlight} 2HCl \] | Non-metal + non-metal |
| 14 | Sodium + chlorine | \[ 2Na + Cl_2 \xrightarrow{\Delta} 2NaCl \] | Ionic compound formation |
| 15 | Cinnabar roasting | \[ 2HgS + 3O_2 \xrightarrow{\Delta} 2HgO + 2SO_2 \] | Roasting (low-reactive metal) |
| 16 | Mercury oxide decomposition | \[ 2HgO \xrightarrow{\Delta} 2Hg + O_2 \] | Thermal decomposition |
| 17 | Roasting zinc blende | \[ 2ZnS + 3O_2 \xrightarrow{\text{roasting}} 2ZnO + 2SO_2 \] | Sulphide ore → oxide |
| 18 | Calcination of zinc carbonate | \[ ZnCO_3 \xrightarrow{\text{calcination}} ZnO + CO_2 \] | Carbonate ore → oxide |
| 19 | Reduction of ZnO by carbon | \[ ZnO + C \xrightarrow{\Delta} Zn + CO \] | Reduction (medium reactive) |
| 20 | Reduction of Fe₂O₃ by CO | \[ Fe_2O_3 + 3CO \xrightarrow{\Delta} 2Fe + 3CO_2 \] | Blast furnace reduction |
| 21 | Molten NaCl electrolysis | \[ 2NaCl(l) \xrightarrow{electricity} 2Na + Cl_2 \] | High-reactive metal (electrolysis) |
| 22 | Alumina electrolysis | \[ 2Al_2O_3(l) \xrightarrow{electricity} 4Al + 3O_2 \] | Aluminium extraction |
| 23 | Rusting of iron | \[ 4Fe + 3O_2 + 2xH_2O \rightarrow 2Fe_2O_3 \cdot xH_2O \] | Corrosion (needs O₂ + H₂O) |
15. Important Definitions (Glossary)
| Term | Definition (board-exam ready) |
|---|---|
| Metal | An element that is usually lustrous, malleable, ductile, sonorous and a good conductor of heat and electricity. |
| Non-metal | An element that is generally dull, brittle, non-sonorous and a poor conductor of heat and electricity. |
| Malleability | The property of a metal to be beaten into thin sheets without breaking. |
| Ductility | The property of a metal to be drawn into thin wires. |
| Sonority | The property of a material to produce a ringing sound when struck; metals are sonorous. |
| Reactivity series | The arrangement of metals in decreasing order of their reactivity. |
| Displacement reaction | A reaction in which a more reactive metal displaces a less reactive metal from a solution of its salt. |
| Ionic compound | A compound formed by the complete transfer of electrons from a metal to a non-metal, held by electrostatic forces (also called an electrovalent compound). |
| Mineral | A naturally occurring substance in the earth's crust in which a metal is present. |
| Ore | A mineral from which a metal can be extracted economically and profitably. |
| Gangue | The earthy impurities (sand, rock) that accompany the ore. |
| Metallurgy | The entire sequence of processes used to extract a metal from its ore (concentration, extraction/reduction and refining). |
| Roasting | Heating an ore (usually a sulphide) in a plentiful supply of air to form the metal oxide. |
| Calcination | Heating an ore (usually a carbonate or hydroxide) in a limited supply of air (or its absence) to form the metal oxide. |
| Reduction | Removal of oxygen (or gain of electrons); the process by which the metal oxide is converted to the metal. |
| Refining | The process of removing impurities from the extracted (impure) metal to obtain the pure metal, e.g. electrolytic refining. |
| Corrosion | The gradual deterioration of a metal due to reaction with air, moisture, acids or other substances in its environment. |
| Alloy | A homogeneous mixture (solid solution) of a metal with one or more other metals or a non-metal such as carbon. |
16. Comparison Tables
A. Metals vs Non-metals
| Property | Metals | Non-metals |
|---|---|---|
| Appearance | Lustrous | Dull (iodine has a slight sheen) |
| Malleability / ductility | Malleable and ductile | Brittle, non-ductile |
| Conduction | Good conductors of heat and electricity | Poor conductors (graphite is an exception) |
| Sonorous | Yes | No |
| State at room temperature | Solid (Hg liquid) | Gas, liquid (Br) or solid |
| Melting / boiling point | Generally high | Generally low |
| Oxides | Basic (amphoteric Al₂O₃, ZnO) | Acidic |
| Ionic behaviour | Form cations (lose electrons) | Form anions (gain electrons) |
| Examples | Fe, Cu, Na, Al, Au, Hg | O, N, C, S, P, Cl, H |
B. Mineral vs Ore
| Basis | Mineral | Ore |
|---|---|---|
| Meaning | Natural source of a metal in the crust | A mineral from which the metal is extracted profitably |
| All ores are minerals? | — | Yes (the reverse is not true) |
| Economy decides? | No | Yes |
| Example | Clay (aluminium) | Bauxite (aluminium) |
C. Roasting vs Calcination
| Basis | Roasting | Calcination |
|---|---|---|
| Ore type | Sulphide ores (sometimes native) | Carbonate / hydroxide ores |
| Air supply | Plentiful (excess oxygen) | Limited (or no) air |
| Purpose | Convert ore to oxide; SO₂ driven off | Drive off H₂O / CO₂ and make oxide |
| Example | 2ZnS + 3O₂ → 2ZnO + 2SO₂ | ZnCO₃ → ZnO + CO₂ |
D. Ionic compounds vs Covalent substances
| Property | Ionic compounds | Covalent substances |
|---|---|---|
| Bonding | Complete transfer of electrons | Sharing of electrons |
| State | Solid at room temperature | Mostly gases / liquids / soft solids |
| Melting / boiling point | High | Generally low |
| Electrical conductivity | Conduct in molten/aqueous state | Usually do not conduct (non-electrolytes) |
| Solubility | Soluble in water, insoluble in organic solvents | Often soluble in organic solvents |
| Examples | NaCl, MgO, CaCl₂ | Water, methane, carbon dioxide |
E. More reactive vs Less reactive metals
| Property | More reactive (K, Na, Ca) | Less reactive (Cu, Ag, Au) |
|---|---|---|
| Electron loss | Lose electrons very easily | Lose electrons with difficulty |
| Reaction with water | React vigorously with cold water | No reaction with water |
| Reaction with dilute acids | Violent — liberate H₂ | No reaction (below hydrogen) |
| Occurrence | Never free; as compounds | May occur free (native) |
| Extraction method | Electrolysis of molten compounds | Roasting / simple heating; no electrolysis |
F. Corrosion vs Prevention methods
| Metal / problem | Form of corrosion | Best prevention |
|---|---|---|
| Iron | Rusting (brown, flaky Fe₂O₃·xH₂O) | Galvanisation, painting, alloying (stainless steel) |
| Copper | Green basic copper carbonate coating | Oiling, lacquering, keeping dry |
| Silver | Black silver sulphide (Ag₂S) | Keeping away from sulphur-containing air; storing in anti-tarnish cloth |
| Machinery / chains | Surface oxidation and wear | Oiling and greasing |
17. Applications & Competency Contexts
Q1. Why is aluminium used to make cooking utensils even though it is reactive with oxygen?
Explanation. Aluminium forms a thin, tough, transparent layer of aluminium oxide (Al₂O₃) on exposure to air. This oxide layer sticks firmly to the surface and prevents further corrosion of the metal underneath — that is why aluminium utensils stay usable. The layer also protects the metal from reacting with food acids, which is why we can cook food in them safely.
Q2. Two metals A and B are compared: A reacts with cold water whereas B reacts only with steam. Which is more reactive, and why?
Explanation. A is more reactive. Reaction with cold water needs a metal to lose electrons readily; steam (hotter) is needed for B, so B loses electrons less easily. In the reactivity series, K/Na/Ca react with cold water while Mg, Al, Zn and Fe need steam — so A is high in the series and B is lower (but still above hydrogen, since it reacts with steam).
Q3. A shop-owner dips iron nails in copper sulphate solution. What would a student observe, and what equation explains it?
Explanation. The blue copper sulphate solution slowly becomes light green (iron(II) sulphate) and a red-brown coating of copper deposits on the nails. Equation: Fe + CuSO₄ → FeSO₄ + Cu. Iron is above copper in the reactivity series, so it displaces copper from its salt solution.
Q4. Why do galvanised iron (GI) water pipes last longer than plain iron pipes?
Explanation. GI pipes are coated with a layer of zinc. Zinc is more reactive than iron, so if the coating is scratched, zinc corrodes first and protects the iron (sacrificial protection); the zinc layer also blocks damp air from the iron. Because the iron does not rust, the pipes last much longer.
Q5. A ring labelled "24 carat gold" bends; one labelled "22 carat" is harder. Explain.
Explanation. Pure (24-carat) gold is soft and bends easily, so jewellers add a small quantity of copper or silver to make 22-carat gold. The alloy is harder, more durable and still lustrous, while costing less gold — alloying improves the hardness of the pure metal.
Q6. Sodium pieces are stored in kerosene, whereas copper utensils are kept in open air. Justify.
Explanation. Sodium reacts vigorously (even explosively) with moisture and oxygen in the air, so it is stored under kerosene, which keeps air and water away. Copper is a low-reactive metal that does not react with air and water at ordinary temperature (it only darkens slowly over years), so it may be kept exposed.
Q7. Electrical wires use copper, but transmission tower cables use an aluminium alloy. Why the difference?
Explanation. Copper is a good conductor and ductile, ideal for house wiring. For overhead transmission lines, a hard aluminium alloy (which is light, strong and cheaper) is preferred to support long spans and reduce sagging. The choice depends on cost, weight and strength, not just conductivity.
Q8. Why is a bridge painted regularly even though it is made of strong steel?
Explanation. Steel is mostly iron, which rusts when air and moisture reach it. Paint forms a barrier between the steel surface and damp air, so the iron does not oxidise. Repainting keeps the barrier intact and prevents the bridge from losing strength to rusting.
18. Competency & Reasoning Questions
Q1. "Sodium is a metal, yet it is so soft that it can be cut with a knife." Do you agree? Give reason.
Reason. Yes. Most metals are hard, but metals have exceptions. Sodium is a soft metal; it has a loose electron cloud and weak interatomic bonding, so it can be cut easily by a knife. Hardness is a general (not universal) metallic property.
Q2. Solid sodium chloride does not conduct electricity, but its aqueous solution does. Why?
Reason. In solid NaCl the Na⁺ and Cl⁻ ions are locked in the crystal lattice and cannot move, so no charge flows. When NaCl dissolves in water (or melts), the lattice breaks apart and the ions become free to move, so the solution conducts electricity.
Q3. A metal M forms oxide MO₂ when burnt in air; the oxide neutralises HCl as well as NaOH. Identify the metal.
Reason. An oxide that neutralises both an acid and a base is an amphoteric oxide. M is aluminium (Al forms Al₂O₃) or zinc (Zn forms ZnO). For MO₂ as formula the safest answer is zinc, whose oxide ZnO neutralises both dilute HCl and NaOH.
Q4. Copper coin placed in silver nitrate solution turns silvery. Explain with the reaction.
Reason. Copper is above silver in the reactivity series, so it displaces silver from silver nitrate. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. Silver metal deposits on the coin, giving it a silvery look, and the solution turns blue from Cu²⁺ ions.
Q5. Why is aluminium extracted by electrolysis but copper is not extracted by electrolysis?
Reason. Aluminium is very reactive; it cannot be reduced by carbon, so the molten oxide is decomposed by electrolysis. Copper is a medium/low reactive metal whose oxide can be reduced by simple heating (roasting), so electrolysis is unnecessary and uneconomical for copper.
Q6. A student saw that a zinc plate in ferrous sulphate solution became covered with a grey layer. Write the reaction and state which metal is safer from rusting afterwards.
Reason. Zinc is above iron, so Zn + FeSO₄ → ZnSO₄ + Fe. Iron deposits on the zinc plate. Zinc, being more reactive, will oxidise in preference to deposited iron — so iron gets protected (sacrificial protection), the same idea as galvanising.
19. Case-Based Questions
Passage 1. A railway bridge built of iron girders develops brown flaky patches when exposed to rain and sea breeze. A maintenance engineer recommends repainting with a zinc-based paint over the cleaned surface.
(a) Name the phenomenon and the two conditions that cause it.
Ans. The phenomenon is corrosion (rusting). It needs both air (oxygen) and moisture (water) to act on the iron.
(b) Give the chemical name and equation for the rust that forms.
Ans. Hydrated iron(III) oxide. 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O.
(c) Why does painting (and using zinc) stop further damage?
Ans. Paint blocks air and moisture from reaching the iron. If the paint is scratched, the zinc layer protects the iron because zinc is more reactive and corrodes first (sacrificial protection). The iron underneath stays intact.
Passage 2. A factory coats iron water tanks with zinc (galvanisation). A worker accidentally scratches the coating on one tank and exposes the iron.
(a) What is the purpose of the zinc coating?
Ans. Zinc forms a protective barrier and, being more reactive than iron, corrodes in place of the iron, preventing rusting.
(b) Even after the scratch, why does the tank remain protected for a while?
Ans. Because zinc is above iron in the reactivity series, it loses electrons (oxidises) more easily than iron. As long as the zinc corrodes, the exposed iron remains spared (sacrificial anode effect). Painting, in contrast, would fail once scratched.
(c) Suggest a cheaper alternative coating and a metal that is not suitable for this purpose.
Ans. Painting or oiling are cheaper alternatives. Copper would not be suitable because copper is less reactive than iron, so it cannot protect iron from rusting.
Passage 3. A student puts a clean copper wire in a test tube of silver nitrate solution and another copper wire in a test tube of copper sulphate solution that already has zinc pieces.
(a) Write the reaction that happens in the silver nitrate tube and the colour change.
Ans. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. The solution turns blue and shiny silver deposits form on the wire because copper is more reactive than silver.
(b) What happens to the zinc pieces in the second tube? Give the equation.
Ans. Zinc displaces copper from copper sulphate: Zn + CuSO₄ → ZnSO₄ + Cu. The blue colour fades and copper deposits, because zinc is more reactive than copper.
(c) Arrange Cu, Ag and Zn in increasing order of reactivity.
Ans. Ag < Cu < Zn.
Passage 4. A metallurgist screens three ores — cinnabar (HgS), zinc blende (ZnS) and bauxite (Al₂O₃·2H₂O) — for extraction by a junior engineer.
(a) Which ore needs only roasting to give the metal directly? Write the equation.
Ans. Cinnabar (HgS) needs only roasting: 2HgS + 3O₂ → 2HgO + 2SO₂, then 2HgO → 2Hg + O₂ (strong heating). Mercury is a low-reactive metal.
(b) Describe the full extraction of zinc from zinc blende.
Ans. (1) Roasting: 2ZnS + 3O₂ → 2ZnO + 2SO₂. (2) Reduction with carbon: ZnO + C → Zn + CO. Zinc is a medium-reactive metal.
(c) Why is bauxite not reduced the same way?
Ans. Aluminium is a high-reactive metal. Its oxide cannot be reduced by carbon; the molten alumina is broken down by electrolysis (2Al₂O₃ → 4Al + 3O₂).
Passage 5. An electrician chooses materials for a new house: copper wire for the room circuit, an aluminium alloy for the overhead service line, and a thin coating option for the outdoor iron gate.
(a) State two physical properties of copper that make it fit for the room circuit.
Ans. Copper is a very good conductor of electricity and is ductile (can be drawn into thin flexible wires) without breaking.
(b) Why is an aluminium alloy used for the overhead line?
Ans. The alloy is light and strong with good conductivity, so long spans between poles can be supported without sagging, and it resists corrosion better than pure aluminium.
(c) The gate is galvanised. Name the coating metal and state how it protects the iron even when scratched.
Ans. Zinc. If the zinc coating is scratched, zinc (more reactive than iron) oxidises first, protecting the iron from rusting.
(d) Give one reason why galvanising is preferred to mere painting for the gate.
Ans. Painting only shields the surface; once a scratch appears, air and moisture reach the iron and rust starts. Galvanising provides sacrificial protection, so the iron stays safe even at scratches.
20. Solved Examples
These 25 examples are arranged: Q1–Q5 basic (1–2 marks), Q6–Q12 moderate (2–3 marks), Q13–Q17 application (3 marks), Q18–Q21 competency (3 marks), Q22–Q25 case-based (3 marks each).
Q1. Define the reactivity series of metals.
Sol. The reactivity series is the arrangement of metals in decreasing order of their reactivity (the ease of losing electrons and undergoing reaction with water, dilute acids, salt solutions, etc.). Hydrogen is included for comparison. The series is K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au.
Q2. What happens when aluminium is heated in air? Write the balanced equation.
Sol. Aluminium reacts vigorously with oxygen to form aluminium oxide, a white solid. Equation: 4Al + 3O₂ → 2Al₂O₃. The oxide is amphoteric (reacts with both acids and bases).
Q3. Define roasting and give one example.
Sol. Roasting is the process of heating a sulphide ore in a plentiful supply of air (excess oxygen) to convert it into the metal oxide and drive off sulphur dioxide. Example: 2ZnS + 3O₂ → 2ZnO + 2SO₂.
Q4. What is galvanisation? Give one use.
Sol. Galvanisation is the process of coating iron or steel with a protective layer of zinc metal to prevent rusting. Use: galvanised iron sheets, pipes and buckets used in damp or outdoor conditions.
Q5. Which metal is a liquid at room temperature, and which non-metal is a liquid at room temperature?
Sol. Mercury (Hg) is the metal that is liquid at room temperature. Bromine (Br₂) is the non-metal that is a liquid at room temperature. (Gallium melts at about 29.8°C, but mercury is the classic answer for a metal liquid at room temperature.)
Q6. Write the balanced equation for the reaction of iron with dilute hydrochloric acid. What is the gas evolved and how is it tested?
Sol. Fe + 2HCl → FeCl₂ + H₂. Hydrogen gas is evolved. The gas is tested by bringing a burning splint near the mouth of the test tube — it burns with a 'pop' sound, confirming H₂.
Q7. Which is more reactive — zinc or copper? State two pieces of evidence from the reactivity series.
Sol. Zinc is more reactive than copper. Evidence: (i) Zinc is placed above copper in the reactivity series; (ii) Zinc reacts with dilute HCl and H₂SO₄ to liberate H₂ gas, while copper does not; (iii) Zinc displaces copper from CuSO₄ solution (Fe + CuSO₄ → FeSO₄ + Cu), but copper cannot displace zinc from ZnSO₄.
Q8. Define calcination and give one example.
Sol. Calcination is the process of heating a carbonate or hydroxide ore in a limited supply of air (or in its absence) to drive off carbon dioxide or water and convert it into the metal oxide. Example: ZnCO₃ → ZnO + CO₂.
Q9. Describe how sodium chloride is formed from sodium and chlorine, and state the type of bond.
Sol. Sodium (metal) loses one electron to form Na⁺: Na → Na⁺ + e⁻. Chlorine (non-metal) gains that electron to form Cl⁻: Cl + e⁻ → Cl⁻. The oppositely charged ions are held by electrostatic attraction, forming the ionic compound sodium chloride, NaCl. The bond type is ionic (electrovalent) bond.
Q10. What is anode mud? Where is it found and what does it contain?
Sol. Anode mud (or anode slime) is a mixture of insoluble impurities that settle below the anode during electrolytic refining of copper. It contains silver (Ag), gold (Au), platinum (Pt) and other non-reactive impurities. The presence of these valuable metals helps recover their cost.
Q11. Name the electrolyte used in the electrolytic refining of copper.
Sol. The electrolyte is acidified copper sulphate solution (CuSO₄ dissolved in water with a small amount of dilute H₂SO₄ added).
Q12. Distinguish between a mineral and an ore with examples.
Sol. A mineral is a naturally occurring substance in the earth's crust in which a metal is present (e.g. clay, an aluminium mineral). An ore is a mineral from which a metal can be extracted economically and profitably (e.g. bauxite, the aluminium ore). All ores are minerals, but not all minerals are ores (clay is a mineral but not an economical ore of aluminium).
Q13. Aluminium is used to make cooking vessels even though it is reactive. Explain.
Sol. On exposure to air, aluminium forms a thin, tough, transparent layer of aluminium oxide (Al₂O₃) on its surface. This oxide layer is non-reactive, sticks firmly and prevents further reaction of the metal with food acids and air. Hence aluminium vessels are safe and durable for cooking.
Q14. A farmer's iron ploughs are rusting. Suggest two methods to protect them and explain why one method gives longer life.
Sol. (i) Painting the plough regularly — paint blocks air and moisture. (ii) Galvanising the plough with zinc — zinc is more reactive than iron and protects it even when scratched (sacrificial protection). Galvanising gives longer life because painting fails at every scratch, whereas zinc continues to protect the iron even when the coating is damaged.
Q15. A student places a zinc strip in iron(II) sulphate solution and a copper strip in zinc sulphate solution. Explain what she observes in each case.
Sol. (a) Zinc + FeSO₄: Zn + FeSO₄ → ZnSO₄ + Fe. Zinc displaces iron because Zn is above Fe. The green colour of FeSO₄ fades and grey iron deposits. (b) Cu + ZnSO₄: No reaction. Copper is below zinc in the reactivity series, so it cannot displace zinc.
Q16. Why are copper cooking vessels no longer common despite copper being an excellent conductor?
Sol. Copper reacts slowly with food acids and acidic substances, forming toxic copper salts. This makes it unsuitable for cooking acidic foods. Stainless steel and aluminium are preferred because they are less reactive with food acids and more hygienic.
Q17. Explain why overhead power transmission lines are made of aluminium alloy rather than copper.
Sol. Although copper is a better conductor, aluminium alloy is much lighter and cheaper, and has sufficient conductivity. Long transmission spans need to be light to prevent sagging and must be strong and corrosion-resistant — an aluminium alloy satisfies all these requirements, making it the practical choice for overhead lines.
Q18. A metal X forms an oxide XO that reacts with hydrochloric acid as well as with sodium hydroxide solution. Identify X and the type of oxide.
Sol. An oxide that reacts with both an acid and a base is an amphoteric oxide. X is zinc (Zn forms ZnO) or aluminium (Al₂O₃; the formula XO corresponds to zinc). X is Zn. Amphoteric oxides: ZnO + 2HCl → ZnCl₂ + H₂O; ZnO + 2NaOH → Na₂ZnO₂ + H₂O.
Q19. Three metals P, Q and R are tested. P displaces Q from its salt solution; Q displaces R from its salt solution; P displaces R from its salt solution. Arrange P, Q and R in increasing order of reactivity.
Sol. P displaces Q → P is more reactive than Q. Q displaces R → Q is more reactive than R. Therefore the increasing order of reactivity is R < Q < P.
Q20. A student mixes sodium hydroxide solution and dilute hydrochloric acid and tests the mixture with litmus paper. State what colour changes she would observe and why.
Sol. NaOH is a base (turns red litmus blue) and HCl is an acid (turns blue litmus red). When they react: NaOH + HCl → NaCl + H₂O. The product is a neutral salt solution. If the amounts are exactly equal, the mixture is neutral and litmus shows no colour change. If excess NaOH remains, the solution stays basic (red litmus turns blue); if excess HCl remains, it stays acidic (blue litmus turns red).
Q21. A silver ornament turns black after a few days in air. Explain the reaction and suggest how the ornament can be protected.
Sol. Silver reacts with sulphur compounds (H₂S) present in the air to form a thin layer of black silver sulphide (Ag₂S): 2Ag + H₂S → Ag₂S + H₂. The black coating is tarnish. Protection: store in an anti-tarnish cloth (which contains a mild sulphur-absorbing agent), or polish regularly with a silver polish, or keep in a sealed bag to limit exposure to air.
Q22. A recycling plant processes old aluminium cans. The cans arrive dirty, are cleaned, melted and reformed.
(a) Give one reason why aluminium is ideal for recycling.
Ans. Aluminium has a low melting point for a metal (about 660°C) and is very ductile, so it can be melted and reformed with relatively low energy compared to extracting aluminium from bauxite by electrolysis, which is extremely energy-intensive.
(b) Explain why the aluminium cans do not rust even when left outdoors in rain.
Ans. Aluminium forms a thin, tough layer of Al₂O₃ on its surface which prevents further reaction. The oxide layer is self-healing — if scratched, it reforms immediately. This is why aluminium resists corrosion.
(c) The recycling plant also receives some iron cans by mistake. How can they be separated from aluminium cans before melting?
Ans. Iron is magnetic; aluminium is not. A magnet can be passed over the incoming cans to pick out the iron cans, separating them from aluminium.
Q23. A ship-builder needs to choose a metal for a seawater tank. Options: iron, aluminium alloy, stainless steel.
(a) Why is plain iron a poor choice for a seawater tank?
Ans. Seawater contains dissolved salts which accelerate rusting. Iron corrodes quickly in salt water, forming hydrated iron oxide (rust), weakening the tank.
(b) Stainless steel is preferred. Name the alloying metals in stainless steel and explain one reason why it resists corrosion.
Ans. Stainless steel contains iron with chromium and nickel (plus carbon). Chromium forms a thin, tough, protective chromium oxide layer on the surface that resists rusting, giving stainless steel its corrosion resistance.
(c) If an aluminium alloy were used instead, explain how it resists corrosion in seawater.
Ans. Aluminium forms a thin Al₂O₃ layer on exposure to air and water, which adheres firmly and protects the metal from further reaction — the same self-healing oxide film that protects aluminium cooking vessels.
Q24. A lab technician extracts copper from copper pyrites (CuFeS₂) by the following steps: roasting in air, then reduction with carbon, then electrolytic refining.
(a) What does roasting achieve? Write one equation.
Ans. Roasting converts the sulphide ore to the oxide and drives off SO₂: 2CuFeS₂ + O₂ → 2Cu₂S + 2FeS + SO₂ (or a simplified board version: 2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂). The goal is to obtain the oxide, which is easier to reduce.
(b) Why is electrolytic refining needed after the reduction step?
Ans. After reduction the copper obtained is impure. Electrolytic refining dissolves impure copper at the anode and deposits pure copper at the cathode (CuSO₄ electrolyte), giving 99.99% pure copper suitable for electrical use.
(c) Name the valuable by-product collected below the anode and explain why it is present.
Ans. Anode mud (containing silver, gold, platinum). These metals are present as impurities in the original ore; they are less reactive than copper, so they do not dissolve at the anode and settle at the bottom as anode mud.
Q25. An iron nail is placed in each of three test tubes: (A) dry air (with a drying agent), (B) oil-covered distilled water (no air contact), (C) ordinary tap water exposed to air. After one week:
(a) In which tube(s) does the nail rust? Why?
Ans. Only tube C rusts. Rusting needs both air (oxygen) and moisture (water) simultaneously. Tube A has air but no water; tube B has water but no air; tube C has both.
(b) Write the equation for the rust formed in tube C.
Ans. 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O (hydrated iron(III) oxide).
(c) A fourth tube (D) contains a nail fully immersed in water with zinc pieces around it. After one week the nail shows no rust. Explain.
Ans. Zinc is more reactive than iron, so the zinc corrodes first (sacrificial protection), leaving the iron uncorroded. This is the principle behind galvanisation.
21. Common Mistakes to Avoid
Mistake 1. Writing the reactivity order incorrectly (placing Fe above Zn or putting H above Cu).
Correction. Memorise K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au. Always double-check the relative position of the two metals you are comparing.
Mistake 2. Saying "metals always react with water".
Correction. Only highly reactive metals (K, Na, Ca) react with cold water; Mg needs hot/steam; Al, Zn, Fe react with steam; Cu, Hg, Ag, Au do not react at all.
Mistake 3. Confusing roasting and calcination.
Correction. Roasting = sulphide ore + excess air (gives oxide and SO₂). Calcination = carbonate/hydroxide ore, limited or no air (drives off CO₂/H₂O).
Mistake 4. Writing "zinc and iron are extracted by electrolysis".
Correction. Only the most reactive metals (above aluminium, plus aluminium itself) are extracted by electrolysis of molten compounds. Zinc/iron are medium-reactive — extracted by reduction of their oxides with carbon/CO.
Mistake 5. Forgetting that rusting needs both air AND moisture.
Correction. Iron in dry air or in water without air does not rust. Rusting happens only when both oxygen and water are present together.
Mistake 6. Writing the corrosion of copper or silver as dark "rust".
Correction. Copper turns green (basic copper carbonate), silver turns black (Ag₂S). Rust specifically refers to hydrated iron(III) oxide on iron.
Mistake 7. Writing the displacement product wrongly by swapping the salt's metal incorrectly.
Correction. In a displacement reaction, the more reactive metal forms its own salt; always keep the accompanying anion (SO₄²⁻, NO₃⁻, Cl⁻) intact: Zn + CuSO₄ → ZnSO₄ + Cu, not ZnCl₂.
Mistake 8. Saying an ionic compound conducts electricity in the solid state.
Correction. Ionic solids have fixed ions that cannot move. They conduct only when molten or dissolved in water, when ions are mobile.
Mistake 9. Believing all metals are hard and all non-metals are gases.
Correction. Sodium and potassium are soft metals; mercury is a liquid metal; bromine is a liquid non-metal; graphite is a non-metal that conducts electricity.
Mistake 10. Writing metal oxides as "acidic".
Correction. Metal oxides are basic (Al₂O₃ and ZnO are amphoteric), while non-metal oxides are acidic (CO₂, SO₂).
Mistake 11. Forgetting anode vs cathode in electrolytic refining.
Correction. Impure copper is the ANODE (it dissolves); pure copper is plated on the CATHODE (negative); electrolyte = acidified copper sulphate.
Mistake 12. Writing "solder is used because it melts at a high temperature so things stick".
Correction. Solder (Pb–Sn) is used precisely because it has a LOW melting point — it melts easily to join wires/parts without damaging them.
Mistake 13. Saying "aluminium is placed directly above magnesium in the reactivity series".
Correction. The order is Mg > Al > Zn > Fe. Between Mg and Al there is no standard metal in the board list, but Al sits below Mg and above Zn; careless ordering here loses easy marks.
22. Board Exam Tips
- Read the reactivity order first. Many 1-mark and 2-mark questions (displacement, extraction method, acid reactivity) are decided purely by the position of the metals — so write the series once on the rough sheet.
- Write balanced equations. Always show states where helpful (s), (aq), (g) and the heat symbol (Δ) when heat is involved. A fully balanced equation is worth more than a word equation.
- Use correct terminology. Do not say "rusting" for copper/silver (that is corrosion/tarnishing); do not say "acidic oxides for metals".
- Answer "why not" question styles. E.g. "Why is copper not extracted by electrolysis?" — answer with the reactivity reason, not just "cheaper".
- For the electrolytic refining diagram, remember it can appear as "draw and label" — practise anode (impure), cathode (pure), electrolyte (acidified CuSO₄) and anode mud once before the exam.
- Case-based questions share the same rules. A 3-mark case passage is answered from the reactivity series and the corrosion/alloying logic — underline the key conditions in the passage before answering.
- Management of time. In a 1-mark question, answer in one line; in a 3-mark question, structure into (i) concept, (ii) equation, (iii) observation/conclusion.
23. Quick Revision
- Physical properties of metals: lustrous, malleable, ductile, sonorous, good conductors; exceptions — Hg liquid, Na/K soft, Ga melts near room temperature.
- Physical properties of non-metals: dull, brittle, non-sonorous, poor conductors; graphite (C) conducts electricity.
- Chemical: metals + O₂ → basic oxides; non-metals + O₂ → acidic oxides; amphoteric oxides (Al₂O₃, ZnO) react with both acids and bases.
- Metal + water: K/Na/Ca react with cold water; Mg, Al, Zn, Fe with steam; Cu onward no reaction.
- Metal + dilute acid: active metals give salt + H₂; pop test confirms H₂.
- Reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au. Higher metal displaces lower metal from its salt — displacement reactions.
- Ionic compounds: formed by electron transfer (metal loses, non-metal gains); hard, high melting, conduct in molten/aqueous state, not as solids.
- Occurrence: mineral (natural source) vs ore (economical mineral); gangue = impurities; low-reactive metals occur native.
- Extraction: low reactivity — heating/roasting (2HgS+3O₂ → 2HgO+2SO₂); medium — roasting/calcination then reduction with C/CO (ZnO+C → Zn+CO; Fe₂O₃+3CO → 2Fe+3CO₂); high — electrolysis of molten compounds (2NaCl → 2Na+Cl₂; 2Al₂O₃ → 4Al+3O₂).
- Refining: electrolytic refining — impure Cu anode, pure Cu cathode, acidified CuSO₄ electrolyte; anode mud has Ag, Au, Pt.
- Corrosion: rusting needs O₂+H₂O; prevention — painting, oiling, galvanising (zinc), alloying, tinning.
- Alloys: brass (Cu+Zn), bronze (Cu+Sn), solder (Pb+Sn, low m.p.), steel (Fe+C), stainless steel (Fe+Cr+Ni, resists rust).
24. Self-Practice Questions (with Answers)
30 questions across 6 types: Conceptual (6), Reactivity & Equations (6), Extraction & Metallurgy (5), Application (5), Competency (5) and Case-based (3). Attempt each before reading the answer.
Conceptual Questions (6)
C1. What is malleability and ductility? Give one use for each property.
Ans. Malleability is the property to be beaten into thin sheets (used to make aluminium foil). Ductility is the property to be drawn into thin wires (used for copper electrical wires).
C2. Name the only two liquid elements at room temperature.
Ans. Mercury (a metal) and bromine (a non-metal).
C3. Why is calcium stored differently from sodium?
Ans. Both react with air/moisture, but sodium reacts far more violently and spontaneously, so it is stored under kerosene. Calcium is less dangerous and is usually kept in airtight containers; both are stored away from air and water.
C4. What is a displacement reaction? Give one example.
Ans. A reaction in which a more reactive element displaces a less reactive element from its compound. Example: Fe + CuSO₄ → FeSO₄ + Cu (iron displaces copper from copper sulphate).
C5. What is an amphoteric oxide? Give two examples.
Ans. An oxide that reacts with both dilute acid and dilute base to give salt and water is an amphoteric oxide. Examples: aluminium oxide (Al₂O₃) and zinc oxide (ZnO).
C6. What is an alloy? Why are alloys preferred over pure metals?
Ans. An alloy is a homogeneous mixture of a metal with another metal or a non-metal. Alloys are preferred because they are harder, stronger, more corrosion-resistant and have tailored properties (e.g. solder's low melting point, stainless steel's rust resistance).
Reactivity & Equations Questions (6)
R1. Write the equation for magnesium reacting with oxygen.
Ans. 2Mg + O₂ → 2MgO (magnesium ribbon burns with a dazzling white flame forming white magnesium oxide).
R2. Write balanced equations for K and for Fe(steam) with water.
Ans. 2K + 2H₂O → 2KOH + H₂. Iron with steam: 3Fe + 4H₂O → Fe₃O₄ + 4H₂.
R3. Arrange Cu, Ag, Zn, Na in increasing order of reactivity.
Ans. Ag < Cu < Zn < Na.
R4. How can you test that the gas evolved when zinc reacts with dilute H₂SO₄ is hydrogen?
Ans. Bring a burning splint to the mouth of the tube — hydrogen burns with a characteristic 'pop' sound. Equation: Zn + H₂SO₄ → ZnSO₄ + H₂.
R5. A copper wire is dipped in a solution of silver nitrate. (a) Write the equation. (b) Name the type of reaction. (c) What colour change occurs?
Ans. (a) Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. (b) Displacement reaction. (c) The solution becomes blue (Cu²⁺ ions) and a grey/white deposit of silver forms on the wire.
R6. Why can a copper strip displace silver, but not zinc, from its salt solution?
Ans. In the reactivity series Cu is above Ag and below Zn. A higher metal displaces a lower one: Cu displaces Ag (Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag), but Cu cannot displace Zn from ZnSO₄ because Cu is below Zn.
Extraction & Metallurgy Questions (5)
E1. What is metallurgy?
Ans. Metallurgy is the entire sequence of processes used to extract a metal from its ore — including concentration of the ore, extraction (reduction) and refining of the metal.
E2. Why is concentration of the ore (removing gangue) the first step of extraction?
Ans. The ore is mixed with gangue (sand, rocky material). Removing gangue enriches the concentration of the metal compound, making the subsequent extraction economical and more efficient.
E3. How is sodium metal obtained from fused NaCl? Write the reaction.
Ans. By electrolysis of molten (fused) sodium chloride: 2NaCl(l) → 2Na(s) + Cl₂(g). Sodium forms at the cathode and chlorine gas at the anode.
E4. Why is the bauxite (alumina) dissolved in molten cryolite before electrolysis?
Ans. Alumina has a very high melting point (about 2050°C). Dissolving it in molten cryolite lowers the working temperature (to about 900–1000°C), saving fuel and making electrolysis economically feasible.
E5. Give the equation for reduction of zinc oxide with carbon.
Ans. ZnO + C → Zn + CO (carbon acts as the reducing agent; heat is applied).
Application Questions (5)
A1. Why are electrical wires made of copper rather than aluminium?
Ans. Copper is an excellent and ductile conductor of electricity and does not oxidise readily in air. Even where aluminium is lighter, for house wiring the superior conductivity and easy joining of copper make it the preferred choice.
A2. An iron gate in a coastal town rusts faster than one inland. Give two reasons.
Ans. (1) Sea breeze carries moisture and salt particles that accelerate rusting. (2) The higher humidity keeps the iron surface always moist. Both factors hasten the corrosion reaction of iron.
A3. Why are cycles and machines greased or oiled?
Ans. Oil or grease creates a protective film that keeps air and moisture away from the metal surface, preventing rusting and reducing friction.
A4. A carpenter nails steel strips with iron nails; after weeks the nails rust. Suggest a better choice and give reason.
Ans. Use galvanised (zinc-coated) iron nails. Zinc protects the iron even when the coating is scratched (zinc corrodes first, safeguarding iron). Plain iron nails rust because they oxidise in air and moisture.
A5. Why does a gold ornament not tarnish, while a silver one does?
Ans. Gold is a noble, least-reactive metal; it does not react with oxygen, moisture or most common substances, so it stays lustrous. Silver reacts with sulphur-containing compounds in air (H₂S) forming black silver sulphide — tarnish.
Competency Questions (5)
K1. A metal X reacts with dilute HCl evolving a gas that burns with a pop. The same metal does not react with cold water but does react with steam. Where would you place X in the reactivity series?
Ans. X reacts with dilute acid (so it is above H) but not with cold water (so it is below Ca). It reacts with steam, so X lies in the Mg-to-Pb part of the series — likely zinc or iron or lead; most likely answers are Zn or Fe.
K2. A solution of copper sulphate is blue. Zinc granules are added and the blue colour disappears. What does the change prove about the relative reactivity of Zn and Cu?
Ans. The blue colour of Cu²⁺ ions disappears because copper(II) has been displaced from the solution by zinc: Zn + CuSO₄ → ZnSO₄ + Cu. This proves zinc (which formed colourless ZnSO₄) is more reactive than copper.
K3. Two powders, A (grey metal) and B (yellow non-metal), are heated together. A flame is produced and a white solid forms. Identify likely A, B and the type of reaction.
Ans. A could be Mg (grey metal) or Al; B is sulphur (yellow non-metal). The white solid is the metal sulphide, e.g. Mg + S → MgS. This is a combination reaction (metal + non-metal forming a compound; ionic for Mg/Al with S).
K4. In the electrolytic refining of copper, the electrolyte does not change in concentration during the process. Explain.
Ans. At the anode, Cu → Cu²⁺ + 2e⁻ (copper dissolves into the solution). At the cathode, Cu²⁺ + 2e⁻ → Cu (copper deposits out). An equal number of Cu²⁺ ions enter and leave the solution, so the CuSO₄ electrolyte concentration remains essentially constant.
K5. A strip of aluminium dipped in copper(II) chloride solution quickly dissolves. However, in dilute hydrochloric acid the same aluminium reacts slowly at first. Why the difference?
Ans. Aluminium needs to shed its protective Al₂O₃ oxide layer before it can react. In copper(II) chloride, chloride ions readily attack and strip this oxide layer quickly, allowing the faster reaction (and copper deposits). In dilute HCl the layer is removed slowly at first; once the layer is disrupted the acid reaction accelerates.
Case-based Questions (3)
Passage (B1). A metal smelting unit owns three ores: haematite (Fe₂O₃), zinc blende (ZnS) and bauxite (Al₂O₃·2H₂O) — but has only one blast furnace with carbon, carbon monoxide and a supply of electricity.
(a) Which ore can be extracted using the blast furnace with carbon/CO? Give the equation.
Ans. Haematite — reduced by CO: Fe₂O₃ + 3CO → 2Fe + 3CO₂. (Zinc blende needs roasting first to ZnO, then reduction with carbon.)
(b) Which ore needs electricity, and why?
Ans. Bauxite — aluminium is a very reactive metal; its oxide cannot be reduced by carbon, so molten alumina is decomposed by electrolysis: 2Al₂O₃ → 4Al + 3O₂.
(c) In what two steps is zinc obtained from zinc blende?
Ans. (1) Roasting: 2ZnS + 3O₂ → 2ZnO + 2SO₂. (2) Reduction: ZnO + C → Zn + CO.
Passage (B2). A museum restorer deals with three ancient objects: a copper statue (green patches), silver coins (black) and an iron sword (brown flakes).
(a) Name the corrosion product on each object.
Ans. Copper: basic copper carbonate (green). Silver: silver sulphide/Ag₂S (black). Iron: hydrated iron(III) oxide, rust (brown flakes).
(b) Write the equation for the rusting of iron.
Ans. 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O.
(c) Suggest how the copper statue could be protected in a humid hall.
Ans. Coat the statue with a transparent protective lacquer (or keep it in a dry, climate-controlled case). This keeps moisture and air away from the metal surface, preventing the formation of the green carbonate coating.
Passage (B3). A railway signal-post is being built. The engineer can choose between plain steel, galvanised steel and stainless steel.
(a) Why do plain-steel posts need frequent repainting?
Ans. Plain steel is mostly iron, which rusts when both air and moisture reach it. Without protection, rusting weakens the metal, so repainting must keep the surface sealed.
(b) Explain how galvanised steel gives longer-lasting protection.
Ans. The zinc layer both keeps air/moisture from reaching the steel and, being more reactive than iron, corrodes preferentially to protect the steel — even at scratches (sacrificial protection). So galvanised posts need far less maintenance.
(c) Give one advantage and one drawback of choosing stainless steel instead.
Ans. Advantage: stainless steel is highly corrosion-resistant and needs no coating — long service life with little maintenance. Drawback: stainless steel is much more expensive than galvanised steel, so costs are higher for a long post or a large structure.
(d) Name the metals alloyed with iron in stainless steel.
Ans. Chromium and nickel (with a little carbon).