A: Metals typically exhibit lustre, are malleable and ductile, conduct heat and electricity, and are sonorous. Examples include iron, copper and aluminium. Non‑metals are mostly dull, brittle (if solid), poor conductors of heat and electricity, and non‑sonorous; examples are sulfur, carbon (graphite) and oxygen. Chemically, metals tend to lose electrons to form positive ions, while non‑metals tend to gain electrons to form negative ions. These differences explain their distinct uses—metals in wiring, construction; non‑metals in insulation, respiration and fertilizers.

A: An alloy is a homogeneous mixture of two or more elements, where at least one is a metal, created to improve strength, corrosion resistance or other properties. Steel (iron + carbon) is harder and stronger than pure iron and is used in construction, tools and vehicles. Brass (copper + zinc) is malleable, corrosion‑resistant and used for musical instruments, decorative items and fittings. Alloys often combine desired properties of different elements for practical applications.

A: Corrosion is the gradual deterioration of metals as a result of chemical reactions with their environment, such as oxidation. Rusting of iron is a common example, producing hydrated iron(III) oxide which weakens structures. Corrosion leads to loss of strength, failure of parts, economic damage and safety hazards—bridges, vehicles and pipelines can fail if not protected. Preventing corrosion is crucial to extend the life of metal objects and save resources.

A: An ore is a rock from which metal compounds can be economically extracted. For iron, the common ore is haematite (Fe₂O₃) or magnetite (Fe₃O₄). Extraction involves concentration (removing impurities), reduction (removing oxygen using carbon in a blast furnace), and purification. In the blast furnace, iron oxide is reduced by carbon monoxide to produce molten iron. The molten iron is then converted to steel or other products. The exact steps vary for different metals but follow concentration, extraction and refining stages.

A: Malleability is the ability of a metal to be hammered or rolled into thin sheets—aluminium can be made into foil and gold into thin leaves. Ductility is the ability to be drawn into wires—copper is drawn into electrical wires. Both properties arise from metallic bonding where layers of atoms can slide over each other without breaking the metallic bond. These properties are important in manufacturing and engineering.

A: Native metals are those found in nature in their metallic, uncombined form, often in riverbeds or rock deposits. They require little or no chemical extraction. Examples include gold and silver, which are often found as nuggets or grains. Such metals are typically less reactive, which is why they do not combine readily with other elements in the Earth’s crust.

A: Metals have a lattice of positive ions surrounded by a sea of delocalised electrons. These free electrons can move throughout the metal, carrying electric charge and thermal energy efficiently. When an electric potential is applied, electrons drift and create current. Heat transfers rapidly as electrons transfer kinetic energy across the metal. This explains why metals like copper and aluminium are used in electrical wiring and cookware.

A: Sonorous means producing a ringing sound when struck. Metals that are hard and elastic, like brass and bronze, vibrate and produce clear tones. This property is used in musical instruments such as bells, cymbals and xylophone bars, where metallic components produce resonant sounds.

A: Non‑metals lack delocalised electrons; their electrons are tightly bound within atoms or covalent bonds, so there are no free charge carriers to conduct electricity. As a result, non‑metals such as sulfur and phosphorus are insulators. Exceptions exist—graphite, a form of carbon, conducts electricity due to delocalised electrons within its layered structure.

A: Physical properties like conductivity, malleability and strength determine use: copper’s excellent conductivity makes it ideal for electrical wiring; aluminium’s low density and strength are useful in aircraft and utensils; iron and steel’s strength and toughness make them suitable for construction. Malleability allows making sheets and foils; ductility enables wires; sonorous metals become musical instruments. Thus, choosing a metal depends on matching its properties to the application.

A: Magnetism in metals arises from unpaired electron spins and the alignment of magnetic domains. Iron, cobalt and nickel have atomic structures that allow magnetic moments to align, creating a net magnetic field. In other metals, electron spins cancel out, and domains do not align, so they are not magnetic. Magnetism is therefore a result of atomic structure and electron configuration.

A: Graphite has a layered structure where each carbon atom is bonded to three others, leaving one delocalised electron per atom. These delocalised electrons can move within the layers, allowing electrical conduction parallel to the layers. Most non‑metals have no delocalised electrons available for conduction, so they act as insulators. Graphite’s unique bonding gives it metallic-like conductivity despite being a non‑metal.

A: Metals react with oxygen to form metal oxides. Reactive metals like magnesium burn in air with a bright flame to form magnesium oxide (2Mg + O₂ → 2MgO). Less reactive metals such as iron form oxides slowly (rusting). The products are basic oxides which can react with acids. The tendency to oxidise depends on the reactivity of the metal.

A: Very reactive metals like potassium and sodium react vigorously with water to form hydroxides and hydrogen gas, often producing heat and flames. Metals like calcium react with water less violently. Metals such as copper and silver do not react with cold water. Based on reactivity, metals are classified as highly reactive (alkali metals), moderately reactive (calcium, magnesium), and least reactive (copper, silver, gold). The reactivity depends on how easily they lose electrons.

A: Many metals react with dilute acids to form a salt and liberate hydrogen gas because metals displace hydrogen from acids. For example: Zn + 2HCl → ZnCl₂ + H₂. The reaction rate depends on the metal’s reactivity; highly reactive metals react readily whereas less reactive metals like copper do not react with dilute acids under normal conditions.

A: Metals typically form basic oxides (e.g., MgO, CaO) which can react with acids to form salts. Non‑metals form acidic or neutral oxides (e.g., CO₂, SO₂) which may form acids when dissolved in water. The nature of the oxide reflects the element’s tendency to lose or gain electrons and its position in the periodic table.

A: Reactive metals such as sodium and potassium react vigorously with air and moisture, sometimes producing flammable hydrogen gas. Storing them under oil prevents contact with air and water, avoiding dangerous reactions. The oil acts as a barrier protecting the metal from oxidation and moisture.

A: A displacement reaction occurs when a more reactive metal displaces a less reactive metal from a compound. For example, iron placed in copper sulfate solution reacts: Fe + CuSO₄ → FeSO₄ + Cu. Iron is more reactive than copper and displaces it from the solution, forming iron sulfate and depositing copper metal.

A: Sulfur reacts with oxygen to form sulfur dioxide: S + O₂ → SO₂. Sulfur dioxide is a gas with a pungent smell and dissolves in water to form sulfurous acid (H₂SO₃), demonstrating how non‑metal oxides can be acidic.

A: The reactivity series ranks metals by their tendency to lose electrons. Metals higher up (e.g., sodium, potassium) react more vigorously with water and acids and can displace metals lower down from compounds. Metals lower in the series (e.g., copper, gold) are less reactive and resist corrosion and reactions. This ranking helps predict reaction outcomes and extraction methods.

A: Rusting is the corrosion of iron in the presence of oxygen and water, producing hydrated iron(III) oxide. Electrolytes like salt accelerate the process by enhancing electron transfer. The overall process involves oxidation of iron to Fe²⁺ and Fe³⁺ ions and subsequent formation of hydrated iron oxides. Rusting weakens iron structures and is a common problem in humid or salty environments.

A: (1) Painting: provides a protective barrier—used on bridges and ships. (2) Oiling/greasing: prevents moisture contact—used on machine parts. (3) Galvanisation: coating with zinc—used for roofing sheets and gutters. (4) Alloying: making stainless steel (iron + chromium) which resists rusting—used in kitchen sinks and cutlery. Each method prevents oxygen/water contact or uses sacrificial protection.

A: Galvanisation coats iron with zinc, which is more reactive and corrodes preferentially. Even if the zinc coating is scratched, the zinc around the exposed area continues to corrode and protects the underlying iron (sacrificial protection). Zinc thereby prevents the formation of iron oxides and extends the life of iron structures and products.

A: Alloying changes the chemical properties of metals, often making them less prone to reaction with the environment. Stainless steel (iron + chromium and sometimes nickel) forms a thin, adherent chromium oxide layer on the surface that prevents further oxidation. This passive layer protects the metal from rusting, making stainless steel suitable for corrosive environments like kitchens and medical instruments.

A: Corrosion leads to structural failures, frequent replacements and maintenance costs across industries—bridges, pipelines and vehicles require repair. Economically, it drains resources and increases safety risks. Environmentally, discarded corroded materials and replacement processes consume energy and raw materials, increasing pollution and resource depletion. Prevention strategies reduce long‑term costs and environmental impact.

A: Selection depends on required properties: conductivity for electrical applications (copper for wires), low density and corrosion resistance for aerospace (aluminium, titanium), hardness and strength for construction (steel). Non‑metals are chosen for insulation and chemical resistance (rubber, plastics, ceramics). Engineers balance properties, cost and environmental factors when choosing materials.

A: Recycling metals reduces the need for mining and ore processing, which are energy‑intensive and environmentally damaging. Recycling aluminium, for example, uses significantly less energy than extracting aluminium from bauxite. Recycled metals also reduce landfill waste and lower greenhouse gas emissions. Thus recycling conserves natural resources and reduces environmental impact.

A: Take three clean iron nails and place one in distilled water, one in tap water, and one in saltwater (salt solution). Leave them exposed to air for several days. Observe that the nail in saltwater rusts faster due to electrolytes that increase conductivity and accelerate corrosion. Record observations and explain that salts act as electrolytes facilitating electron flow and oxidation.

A: Copper has higher electrical conductivity and is more ductile and reliable than aluminium. It forms durable connections with less thermal expansion and fewer failures. Although aluminium is lighter and cheaper, copper’s superior long‑term performance makes it preferred for many wiring applications, especially where durability and safety are priorities.

A: Key points: Metals are lustrous, malleable, ductile, conductive and often sonorous; non‑metals are generally dull, brittle and poor conductors. Metals react with oxygen, water and acids to form oxides, hydroxides and salts; non‑metals form acidic oxides. Rusting of iron is a common corrosion; prevention methods include painting, oiling, galvanisation and alloying. Alloys improve properties (steel, brass). Understand simple reactions, reactivity trends and real‑life uses for good answers in exams.