Metallurgy Class 10 ICSE — Rachna Chapter 7

Metallurgy class 10 ICSE is the chapter that explains how metals are extracted from their ores and turned into everyday materials. This Rachna Sagar Chemistry chapter covers the activity series, corrosion, extraction of aluminium, forms of iron, uses of zinc and alloys. You will learn definitions, equations and the reasoning behind each extraction method. The sections below follow the chapter in the book’s own order.

Download Metallurgy Class 10 ICSE Chemistry Chapter 7 PDF

Download the Metallurgy Class 10 ICSE Chemistry chapter PDF from the Rachna Sagar textbook to revise offline or print for practice. It is the full chapter as printed in the book — 31 pages, 35 sections, 87 questions.

What is inside Chapter 7

31 pages · 35 sections · 87 exercise questions

• EXERCISE-7A: 11 questions
• EXERCISE-7B: 12 questions
• EXERCISE-7C: 18 questions
• EXERCISE-7D: 12 questions
• EXERCISE-7E: 34 questions

7.1 INTRODUCTION

The chapter opens with the distribution of elements in the earth’s crust. Oxygen is the most abundant at 46.6%, followed by silicon at 27.7%. Aluminium (8%) and iron (5%) are next. Of the 118 known elements, only 22 are non-metals. All metals except mercury are solids at room temperature. Bromine is the only liquid non-metal.

Worked pattern: rank aluminium, iron, calcium, sodium, potassium and magnesium by abundance. The order is aluminium, iron, calcium, sodium, potassium, magnesium. This matches decreasing crustal abundance.

7.2 METALS

Metals are defined as elements that form positive ions by losing electrons. Sodium loses one electron to form \mathrm{Na}^{+}. Calcium loses two to form \mathrm{Ca}^{2+}. Aluminium loses three to form \mathrm{Al}^{3+}. The electrons lost equal the metal’s valency. Copper was the first metal used by man. Hydrogen is a non-metal but forms the positive ion \mathrm{H}^{+}.

Worked example: magnesium has configuration 2, 8, 2. It loses two outer electrons to form \mathrm{Mg}^{2+}. Its valency is therefore 2. This matches the pattern of ion formation in this ICSE chapter.

7.3 NON-METALS

Non-metals form negative ions by gaining electrons. Chlorine gains one to form \mathrm{Cl}^{-}. Oxygen gains two to form \mathrm{O}^{2-}. Nitrogen gains three to form \mathrm{N}^{3-}. Electrons gained equal valency. The major metals in the crust by decreasing abundance are aluminium, iron, calcium, sodium, potassium and magnesium. Major non-metals are oxygen, silicon, phosphorus and sulphur.

Worked example: oxygen has six valence electrons. It gains two more to complete its octet. The equation is \mathrm{O} + 2\mathrm{e}^{-} \rightarrow \mathrm{O}^{2-}.

7.4 CATEGORISATION OF METALS

Metals are grouped by their position and shared characteristics in the Periodic Table. Four families matter for this chapter: alkali metals, alkaline earth metals, transition metals and inner transition metals.

7.4.1 Alkali Metals

Lithium, sodium, potassium, rubidium, caesium and francium form group IA. Each has one valence electron and valency one. They react with water to form strong alkalis, which is why they are called alkali metals. They are soft, easily cut with a knife, and very reactive. They react violently with water and dilute acids to give hydrogen. They do not occur free in nature.

7.4.2 Alkaline Earth Metals

Beryllium, magnesium, calcium, strontium, barium and radium form group IIA. Each has two valence electrons and valency two. They are less reactive than alkali metals but still never occur free. They are harder than alkali metals and react with water and acids to produce hydrogen. Their hydroxides are alkalies, and their oxides occur in the earth’s crust.

7.4.3 Transition metals

These sit in groups 3 to 12 of the Periodic Table. Iron, zinc, copper and mercury are examples. They have high melting and boiling points, conduct heat and electricity well, and many act as catalysts. Most show variable valency. Iron can form both \mathrm{Fe}^{2+} and \mathrm{Fe}^{3+}.

7.4.4 Inner transition metals

The lanthanides (period 6) and actinides (period 7) sit at the bottom of the Periodic Table. They are heavy metals, show variable valency and form coloured ions. Many actinides are radioactive. Metalloids such as boron, silicon, germanium, arsenic, antimony, tellurium and polonium show both metallic and non-metallic properties.

7.5 POSITION OF METALS AND NON-METALS IN THE PERIODIC TABLE

Metals occupy the left side and middle of the Periodic Table. Non-metals sit on the right. Hydrogen is the only non-metal placed on the left. A zig-zag line of metalloids separates the two groups.

7.6 COMPARISON OF METALS AND NON-METALS (GENERAL PROPERTIES)

Physically, metals are solid except mercury and gallium, lustrous except lead, and dense except sodium and potassium. Non-metals exist in all three states. Bromine is the liquid non-metal. Iodine and graphite are lustrous. Diamond is the hardest known substance. This comparison builds on concepts from other Class 10 ICSE Chemistry chapters.

Chemically, metals lose 1 to 3 electrons to form cations. They discharge at the cathode during electrolysis. Their oxides are basic, with amphoteric examples like \mathrm{Al}_{2}\mathrm{O}_{3}, \mathrm{PbO} and \mathrm{ZnO}. Non-metals gain electrons to form anions, discharge at the anode, and form acidic oxides. Neutral oxides include CO, NO and \mathrm{N}_{2}\mathrm{O}.

Worked example: classify \mathrm{Fe}_{2}\mathrm{O}_{3}, \mathrm{Mn}_{2}\mathrm{O}_{7}, \mathrm{PbO} and NO. \mathrm{Fe}_{2}\mathrm{O}_{3} is basic. \mathrm{Mn}_{2}\mathrm{O}_{7} is acidic. \mathrm{PbO} is amphoteric. NO is neutral.

7.7 THE ACTIVITY SERIES OF METALS

The activity series arranges metals in decreasing order of reactivity. Potassium is most reactive; gold is least. Hydrogen sits in the series because it forms a positive ion. The full order is K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au.

Memory aid: King Never Comes Marching Alone, Zebras For Peace Here. Can He Get Gold? The first letters match the series order.

Electropositive character decreases down the series. K and Na react vigorously with cold water, producing ignited hydrogen. Mg reacts with steam. Al, Zn and Fe react with steam when hot. Metals above hydrogen displace hydrogen from dilute acids. Lead reacts slowly because its chloride and sulphate coat the metal. Higher metals displace lower metals from salt solutions.

Worked example: place a zinc strip in copper sulphate solution. Zinc is higher in the series. It displaces copper: \mathrm{Zn} + \mathrm{CuSO}_{4} \rightarrow \mathrm{ZnSO}_{4} + \mathrm{Cu}. The blue solution fades as copper deposits.

7.7.1 Action of heat on some metallic compounds

This section summarises what happens to oxides, hydroxides, carbonates and nitrates when heated. Oxides from K to Al are stable to heat and need electrolysis for reduction. Hydroxides of Ca, Mg, Zn, Pb and Cu decompose to the oxide and water. Carbonates decompose to the oxide and carbon dioxide. Nitrates form the oxide, nitrogen dioxide and oxygen. Oxides of Hg and Ag decompose to the metal and oxygen on heating.

7.8 CORROSION

Corrosion is the attack on a metal surface by air, moisture or other substances. Alkali metals form carbonates through a series of reactions. Aluminium and magnesium form white oxides. Iron forms brown rust, \mathrm{Fe}_{2}\mathrm{O}_{3} \cdot x\mathrm{H}_{2}\mathrm{O}. Lead forms basic lead carbonate. Copper forms green basic copper carbonate. Silver tarnishes black as \mathrm{Ag}_{2}\mathrm{S}. Noble metals like gold and platinum do not corrode easily.

7.8.1 Corrosion of Iron

Rusting is the slow oxidation of iron by atmospheric oxygen in the presence of water. The equation is 4\mathrm{Fe} + 3\mathrm{O}_{2} + 2x\mathrm{H}_{2}\mathrm{O} \rightarrow 2\mathrm{Fe}_{2}\mathrm{O}_{3} \cdot x\mathrm{H}_{2}\mathrm{O}. Rust is soft and powdery. It falls off, exposing fresh iron to further rusting. The classic activity proves both air and water are needed. Nails in moist air rust; nails in dry air or boiled water alone do not.

7.8.2 Factors favourable for increased corrosion

Four factors increase corrosion. Metals higher in the activity series corrode more easily. Dissolved salts act as electrolytes and speed corrosion. A less reactive metal in contact with iron increases rusting. Pollutants like \mathrm{NO}_{2} and \mathrm{CO}_{2} also increase rusting.

7.8.3 Prevention of corrosion

Three methods prevent corrosion. Barrier protection coats the metal so air and water cannot reach it. Painting, oiling and electroplating are examples. Sacrificial protection covers iron with a more electropositive metal like zinc. This is galvanisation. Cathode protection connects iron to a more reactive metal by wire. The reactive metal corrodes instead of the iron.

7.8.4 Advantage of corrosion

Corrosion can protect a metal. Aluminium and zinc form sticky, impervious oxide layers. These layers shield the metal underneath from further attack. This is why aluminium articles resist corrosion.

7.9 OCCURRENCE OF METALS

Most metals occur in combined state. Reactive metals like K and Na occur as chlorides. Ca and Mg occur as carbonates. Al occurs as oxide. Zn, Pb, Cu, Hg and Ag occur as sulphides. Gold and platinum occur free. The ores table lists carnallite for K, rock salt for Na, limestone for Ca, magnesite for Mg, bauxite for Al, zinc blende for Zn, haematite for Fe, galena for Pb, copper pyrites for Cu, cinnabar for Hg and argentite for Ag.

Key terms: metallurgy is the extraction of pure metals from ores. Minerals are natural metal compounds mixed with impurities. Gangue is the earthy impurity. Flux removes gangue. Slag is the fusible product of flux and gangue. Smelting reduces oxide ore with flux. Worked pattern: \mathrm{SiO}_{2} gangue reacts with \mathrm{CaCO}_{3} flux to give \mathrm{CaSiO}_{3} slag and \mathrm{CO}_{2}. You can verify the current Class 10 Chemistry syllabus on the official CISCE website.

7.10 STEPS INVOLVED IN EXTRACTION

Extraction begins with pulverisation. Ores are crushed into fine powder in jaw crushers and ball mills. This prepares the ore for concentration.

7.10.2 Dressing (Concentration) of ores

Four methods concentrate ores. Hydrolytic or gravity separation uses density difference. Denser ore settles in grooves; lighter gangue washes away. Magnetic separation uses magnetic properties. Froth floatation uses oil to wet sulphide ores. They float as froth while gangue sinks. Leaching dissolves the ore in a reagent, filters impurities, then precipitates the ore.

7.10.3 Roasting and Calcination

Roasting heats sulphide ores in excess air. Zinc blende gives zinc oxide and \mathrm{SO}_{2}: 2\mathrm{ZnS} + 3\mathrm{O}_{2} \rightarrow 2\mathrm{ZnO} + 2\mathrm{SO}_{2} at 800–900°C. Calcination heats carbonate and hydrated ores in the absence of air. Calamine gives ZnO and \mathrm{CO}_{2}: \mathrm{ZnCO}_{3} \rightarrow \mathrm{ZnO} + \mathrm{CO}_{2}. Both processes make the ore porous and more reactive.

7.10.4 Reduction of metal oxides

Highly reactive metals such as K, Na, Ca, Mg and Al need electrolytic reduction. Their oxides cannot be reduced by carbon, CO or hydrogen. Middle metals like Fe, Zn, Pb and Cu use carbon reduction. Zinc oxide is reduced only by carbon: \mathrm{ZnO} + \mathrm{C} \rightarrow \mathrm{Zn} + \mathrm{CO}. Iron’s haematite is reduced in a blast furnace with coke and limestone. Metals low in the series, like Hg and Ag, need only heating.

Their oxides decompose to the metal and oxygen above 300°C. Manganese oxide is reduced by aluminium powder in an exothermic reaction: 3\mathrm{MnO}_{2} + 4\mathrm{Al} \rightarrow 3\mathrm{Mn} + 2\mathrm{Al}_{2}\mathrm{O}_{3} + \text{Heat}.

Worked example: a metal placed just below hydrogen, like copper, is obtained from its oxide by carbon reduction. \mathrm{CuO} + \mathrm{C} \rightarrow \mathrm{Cu} + \mathrm{CO}.

7.10.5 Refining of the metal

Three methods refine metals. Distillation works for volatile metals like mercury and zinc. Liquation works for low melting metals like lead and tin. Electro-refining uses the impure metal as anode and pure metal as cathode. Copper, aluminium and zinc are refined this way. This extraction route follows the same pattern seen in other ICSE Class 10 books.

7.11 COMMON ORES OF ALUMINIUM, IRON AND ZINC

Aluminium ores: bauxite (\mathrm{Al}_{2}\mathrm{O}_{3} \cdot 2\mathrm{H}_{2}\mathrm{O}), cryolite (\mathrm{Na}_{3}\mathrm{AlF}_{6}), corundum (\mathrm{Al}_{2}\mathrm{O}_{3}). Iron ores: red haematite (\mathrm{Fe}_{2}\mathrm{O}_{3}), brown haematite (2\mathrm{Fe}_{2}\mathrm{O}_{3} \cdot 3\mathrm{H}_{2}\mathrm{O}), magnetite (\mathrm{Fe}_{3}\mathrm{O}_{4}), iron pyrites (\mathrm{FeS}_{2}), siderite (\mathrm{FeCO}_{3}). Zinc ores: zinc blende (\mathrm{ZnS}), zincite (\mathrm{ZnO}), calamine (\mathrm{ZnCO}_{3}).

7.12 ALUMINIUM

Symbol Al, atomic number 13, relative atomic mass 27, valency 3. Electronic configuration 2, 8, 3. Silvery white metal with specific gravity 2.2. Position: period 3, group IIIA (13).

7.12.1 Discovery and Occurrence

Aluminium was discovered by Oersted in 1825. It is the most abundant metal in the earth’s crust. It is reactive, so it never occurs free in nature.

7.12.2 Extraction of Aluminium

Hall and Heroult developed the extraction process in 1885. Bauxite contains 60% \mathrm{Al}_{2}\mathrm{O}_{3}. Baeyer’s process purifies it. Bauxite dissolves in caustic soda to form sodium meta aluminate. Dilution hydrolyses it to precipitate \mathrm{Al}(\mathrm{OH})_{3}. Ignition at 1000°C gives alumina. Alumina melts at about 2050°C. At this temperature aluminium vaporises. Cryolite lowers the fusion temperature to 950°C.

The Hall-Heroult cell uses alumina 20%, cryolite 60% and fluorspar 20%. Cathode reaction: 4\mathrm{Al}^{3+} + 12\mathrm{e}^{-} \rightarrow 4\mathrm{Al}. Anode reaction: 6\mathrm{O}^{2-} - 12\mathrm{e}^{-} \rightarrow 6[katex display="true"]\mathrm{O}[/katex], forming \mathrm{O}_{2}, then CO and \mathrm{CO}_{2}. The graphite anode oxidises and must be replaced. Aluminium obtained is 99.8% pure.

Worked example: to deposit 4 moles of aluminium, 12 moles of electrons are needed. Each \mathrm{Al}^{3+} ion accepts 3 electrons. So 4 \mathrm{Al}^{3+} ions accept 12 electrons total.

7.12.3 Properties of aluminium

Physical: silvery white, light but strong, malleable, ductile, good conductor of heat and electricity. Melting point 660°C, boiling point 2050°C. Chemical: oxidises in air to form a protective \mathrm{Al}_{2}\mathrm{O}_{3} layer. Reacts with steam to give hydrogen. Amphoteric — reacts with both acids and bases.

With NaOH it forms sodium meta aluminate. With dilute HCl it gives \mathrm{AlCl}_{3} and hydrogen. With dilute \mathrm{H}_{2}\mathrm{SO}_{4} it gives \mathrm{Al}_{2}(\mathrm{SO}_{4})_{3} and hydrogen. Concentrated \mathrm{H}_{2}\mathrm{SO}_{4} gives \mathrm{SO}_{2}. Nitric acid renders it passive. It is a powerful reducing agent in aluminothermy.

7.13 USES OF ALUMINIUM

Used in cooking utensils, building construction, packaging foils, furniture and household gadgets. Used as a deoxidizer in steel manufacture, in flash bulbs, explosives like ammonal, reflecting telescope mirrors, power cables and shipbuilding. Used in aluminothermy for welding broken iron girders. Thermite is 3 parts ferric oxide with 1 part aluminium powder. The reaction reaches about 3000°C.

7.14 FORMS AND USES OF IRON

Quenching heats steel above 800°C and plunges it into cold water or oil. This makes it harder and more brittle. Annealing reheats quenched steel to 220–300°C and cools it slowly. This removes brittleness. Mild steel (0.15–0.5% carbon) makes nuts and bolts. Hard steel (0.5–1.5% carbon) makes tools.

7.15 USES OF ZINC

Zinc is mostly used for galvanisation — coating iron and steel sheets to prevent rusting. It makes alloys like brass, bronze and German silver. It forms the negative electrode in dry cells. It is used in lithographic plates and as a reducing agent in organic reactions.

7.16 ALLOYS

An alloy is a homogeneous mixture of two or more metals, or metals with non-metallic elements. Alloy properties differ greatly from the components. Gold is too soft without copper. Stainless steel has 15–18% chromium. Fusible alloys melt between 51°C and 260°C. An amalgam is an alloy of mercury with other metals. Mercury does not amalgamate with iron, so it is stored in iron bottles.

7.16.2 Reasons for alloying

Alloys are made to change specific properties. Aluminium bronze looks like gold (colour). Sodium amalgam is less reactive than sodium (chemical activity). Type metal is easily cast (casting ability). Solder melts at 180°C, lower than lead or tin (melting point). Brass is harder than copper (hardness). Nichrome has more electrical resistance than copper (resistance).

7.16.3 Method of making alloys

Two methods exist. Fusing metals together, as with brass from melting zinc and copper. Or compressing finely divided metals, as with wood metal — a fusible alloy of lead, tin, bismuth and cadmium used in automatic sprinklers.

7.16.4 Alloys as solid solutions

An alloy is a homogeneous mixture, which means it behaves as a solid solution of one metal in another. This explains why alloy properties differ from the pure components that form them.

What to practise from Metallurgy Class 10 ICSE

Work through these exercise blocks from the chapter in order. Each one targets a different part of the syllabus.

• EXERCISE-7A — 11 questions on metals and non-metals
• EXERCISE-7B — 12 questions on the activity series
• EXERCISE-7C — 18 questions on extraction
• EXERCISE-7D — 12 questions on aluminium
• EXERCISE-7E — 34 questions on uses of metals and alloys

In the ICSE board exam, questions from this chapter typically test the activity series order, the extraction method matched to a metal’s position, and alloy compositions and uses.

Metallurgy Class 10 ICSE frequently asked questions

Common mistakes students make: confusing roasting with calcination, assuming aluminium corrodes like iron, and forgetting that gallium and mercury are liquid metals. Roasting uses excess air; calcination uses no air. Aluminium’s oxide layer protects it. Mercury and gallium are the liquid metals.

What is metallurgy in class 10 ICSE chemistry?

Metallurgy is the process used to extract metals in pure form from their ores. It covers production, purification of metals and manufacture of alloys. The term appears in section 7.9 of this chapter.

Why is aluminium extracted by electrolysis and not by carbon reduction?

Alumina has great affinity for oxygen. It is a very stable compound. Carbon, carbon monoxide and hydrogen cannot reduce it. Electrolytic reduction is the only practical method.

Why does aluminium not corrode even though it is more reactive than iron?

Aluminium forms a thin, tough layer of aluminium oxide on its surface. This layer is sticky and impervious. It protects the metal underneath from further attack. Iron’s rust flakes off and exposes fresh metal.

What is the difference between roasting and calcination?

Roasting heats the ore in excess air. It is used for sulphide ores. Calcination heats the ore in the absence of air. It is used for carbonate and hydrated ores.

Which metals are liquid at room temperature?

Mercury is the only metal that is liquid at room temperature. Gallium changes to liquid when handled. Bromine is the liquid non-metal.

Reference: ICSE Class 10 Chemistry textbook material, chapter 7: Metallurgy — CISCE-aligned.

More from this book

• Chapter 3: Study of Acids Bases and Salts
• Chapter 10: Nitric Acid