Chapter 4 — Materials Metals And Non Metals

Materials around us are broadly classified as metals and non-metals. The Class 8 chapter 'Materials: Metals and Non‑metals' begins with an introduction that explains how these two groups differ in physical and chemical properties, how they are obtained and used, and some important exceptions.

What are metals? Metals are elements that usually have a shiny appearance (metallic lustre), are good conductors of heat and electricity, are malleable (can be beaten into sheets), ductile (can be drawn into wires), and are generally solid at room temperature (except mercury). Chemically, many metals tend to lose electrons and form basic oxides. Metals react with acids to produce hydrogen gas and often react with oxygen to form metal oxides.

What are non‑metals? Non‑metals are elements that are generally dull (lack metallic lustre), are poor conductors of heat and electricity (many are insulators), and are brittle if solid (they break or crumble when struck). Non‑metals form acidic or neutral oxides and often gain electrons in chemical reactions. Their physical and chemical behaviour is quite different from metals.

Important points and exceptions:

• Mercury is a metal but is liquid at room temperature.
• Graphite (a form of carbon) is a non‑metal but conducts electricity.
• Sodium and potassium are metals that are soft and can be cut with a knife.

Occurrence and extraction (brief): Metals usually occur in ores (compounds of metals with other elements). These ores are concentrated and then metals are extracted by processes such as heating with carbon or electrolysis, depending on the reactivity of the metal.

Uses: Metals (iron, copper, aluminium, gold) are widely used in construction, electrical wiring, utensils and jewellery. Non‑metals (oxygen, nitrogen, carbon, sulphur, chlorine) are essential in life processes, medicines, fertilizers, bleaching agents and many industrial chemicals.

How to distinguish quickly: Simple tests in the laboratory or at home (with care and appropriate supervision) include testing for lustre, malleability, sonorousness (sound when struck), and electrical conductivity.

Definition: Metals are a class of materials that share a set of characteristic physical properties due to metallic bonding (a lattice of positive ions surrounded by delocalised electrons). These properties make metals useful in everyday life and engineering.

• Lustre (Shiny surface): Metals reflect light from their surfaces, giving a metallic shine. Example: polished steel or aluminium cookware.
• Malleability: Metals can be beaten or rolled into thin sheets without breaking because layers of atoms can slide over each other while remaining bonded. Example: aluminium foil, gold leaf.
• Ductility: Metals can be drawn into wires. This is due to the same sliding of atomic planes. Example: copper wires used in electrical circuits.
• Electrical conductivity: Metals conduct electricity because delocalised electrons can move freely through the metal. Example: copper and aluminium in power cables.
• Thermal conductivity: Metals conduct heat well as moving electrons and lattice vibrations transfer thermal energy. Example: copper bottoms of cooking pans and aluminium radiators.
• High density: Most metals have higher densities compared to non-metals because metal atoms are closely packed. Example: lead and gold are very dense; aluminium is relatively low-density among metals.
• High melting and boiling points: Many metals require large amounts of energy to break metallic bonds, so they often have high melting/boiling points. Example: iron and tungsten.
• Hardness and strength: Metals vary widely in hardness and tensile strength; some (like steel) are very strong and used in construction, while others (like sodium) are soft.
• Sonorous: Metals produce a ringing sound when struck (useful for bells and musical instruments).
• Opaque: Bulk metals are generally not transparent because they reflect/absorb visible light.

Why these properties occur: Metallic bonding gives metals a sea of free electrons. These electrons allow easy flow of charge (electrical conductivity), transfer of thermal energy (thermal conductivity), and allow atomic layers to move without breaking bonds (malleability and ductility). Close packing of atoms leads to high density and strong interatomic forces give high melting points.

Simple classroom observations/experiments: Hammer a small piece of aluminium into a thin sheet (malleability), pull copper wire to show ductility, touch metal spoon and plastic spoon near a flame to see thermal conduction, strike a metal bell to hear sonorous property.

Definition: Non-metals are elements that generally lack the typical physical properties of metals. They show a wide range of physical behaviours because they exist as gases, liquids or solids at room temperature and have different types of bonding (molecular or covalent network).

Main physical properties:

• States: Non-metals are found in all three states at room temperature — many are gases (oxygen, nitrogen, chlorine), a few are liquids (bromine), and several are solids (carbon, sulfur, phosphorus).
• Appearance / Lustre: Most non-metallic solids are dull and lack metallic lustre. Exceptions: iodine shows a shiny crystalline appearance; graphite appears lustrous along cleavage planes.
• Malleability and Ductility: Non-metals are not malleable or ductile. Solid non-metals are usually brittle and break or shatter when hammered (e.g., sulfur, phosphorus).
• Hardness: Varies widely. Many non-metals are soft (sulfur, phosphorus), but some are extremely hard (diamond — an allotrope of carbon is the hardest known natural substance).
• Density: Generally have lower densities than metals. Most non-metals are light (e.g., hydrogen, helium, nitrogen), though there are exceptions.
• Melting and Boiling Points: Most molecular non-metals have low melting and boiling points due to weak intermolecular forces (e.g., oxygen, nitrogen). Covalent-network non-metals (diamond, graphite) have very high melting points.
• Electrical Conductivity: Non-metals are typically poor conductors of electricity. Important exception: graphite (a form of carbon) conducts electricity because of delocalised electrons in its layers.
• Thermal Conductivity: Generally poor thermal conductors (insulators). Diamond is a notable exception — it conducts heat very well.
• Sonority: Non-metals are non-sonorous (do not produce sound on striking), unlike metals which are sonorous.

Why these properties? The physical behaviour of non-metals arises from their bonding and structure. Many non-metals form simple molecules held together by weak intermolecular forces (low melting/boiling points, poor conductivity). Others form giant covalent networks (diamond, graphite) giving high hardness or high melting points and, in graphite’s case, electrical conductivity along layers.

Typical classroom comparisons: Compare a piece of sulphur (soft, brittle, dull, poor conductor) with a piece of graphite (soft, slippery, good conductor) and diamond (extremely hard, transparent), to see the variety among non-metals.

What are chemical properties of metals?
Chemical properties describe how metals react with other substances (oxygen, water, acids, salts) to form new compounds. Metals tend to lose electrons and form positive ions (cations); this makes most metal reactions oxidation–reduction (redox) processes.

• Reaction with oxygen (formation of oxides)
  Most metals react with oxygen to form metal oxides. These oxides are usually basic (e.g., Na2O, CaO). Some metal oxides are amphoteric (react with both acids and bases) — e.g., Al2O3 and ZnO. Example equation: 2Mg + O2 → 2MgO. Rapid oxidation (combustion) produces bright flames for active metals like Mg.
• Reaction with water
  Reactivity with water depends on the metal:
  - Very reactive metals (Group 1: K, Na) react violently with cold water to form hydroxides and hydrogen gas (2Na + 2H2O → 2NaOH + H2).
  - Less reactive metals (Mg) do not react with cold water but react with steam to give oxide and hydrogen (Mg + H2O (steam) → MgO + H2).
  - Metals like Cu, Ag, Au do not react with water.
• Reaction with acids
  Most active metals react with dilute acids to form a salt and hydrogen gas. Hydrogen evolution (test: 'pop' sound) shows H2 formation. Example: Fe + 2HCl → FeCl2 + H2. Metals above hydrogen in the reactivity series displace H+ from acids.
• Displacement reactions (metal + metal salt solution)
  A more reactive metal displaces a less reactive metal from its salt solution. This follows the reactivity series. Example: Zn + CuSO4 → ZnSO4 + Cu (zinc displaces copper).
• Corrosion (rusting of iron)
  Corrosion is slow oxidation in presence of air and moisture. Iron rusts to form hydrated iron(III) oxide (commonly written as Fe2O3·nH2O). Rusting needs O2 and H2O and is promoted by salts and acids. Prevention: painting, oiling, galvanization (zinc coating), sacrificial protection (attach a more reactive metal).
• Amphoteric behaviour
  Some metal oxides (Al2O3, ZnO) react with both acids and bases. Example: ZnO + 2HCl → ZnCl2 + H2O and ZnO + 2NaOH + H2O → Na2[Zn(OH)4] (or simplified: ZnO + 2NaOH → Na2ZnO2 + H2O in basic medium).
• Role of electrons (redox view)
  Metals lose electrons (oxidation). Example: Mg → Mg2+ + 2e− (oxidation). Oxygen or hydrogen ions commonly accept electrons (reduction). This explains why metals reduce H+ to H2 and also why displacement reactions occur.

Reactivity series (useful summary): K > Na > Ca > Mg > Al > Zn > Fe > H > Cu > Ag > Au. Metals above H can displace H from acids; a metal higher in the series displaces a metal lower in the series from solution.

Practical points / experiments: Observe the pop sound when metal reacts with acid (hydrogen test); place an iron nail in copper sulphate to see copper deposition; test sodium or potassium reactions with water (only under supervised conditions); compare corrosion rates with/without protective coatings.

Summary: Chemical properties of metals include reactions with oxygen, water, acids and salts, displacement reactions, corrosion and amphoterism. These depend on the metal’s position in the reactivity series and its tendency to lose electrons.

Overview
Non-metals are elements that typically gain electrons in chemical reactions, form covalent bonds, and show a variety of chemical behaviours such as forming acidic oxides, forming hydrides, reacting with oxygen, and acting as oxidising agents. They usually do not displace hydrogen from acids and do not show metallic character.

Main chemical properties

• Reaction with oxygen (combustion): Most non-metals react with oxygen to form oxides. These are often acidic oxides. Examples: C + O₂ → CO₂; S + O₂ → SO₂; 4P + 5O₂ → P₄O₁₀. These reactions release energy (combustion).
• Formation of acids from oxides: Many non-metal oxides dissolve in water to give acids: CO₂ + H₂O → H₂CO₃ (carbonic acid); SO₂ + H₂O → H₂SO₃ (sulphurous acid); SO₃ + H₂O → H₂SO₄ (sulphuric acid); P₄O₁₀ + 6H₂O → 4H₃PO₄ (phosphoric acid).
• Reaction with hydrogen (hydride formation): Non-metals combine with hydrogen to form covalent hydrides, e.g. 2H₂ + O₂ → 2H₂O; C + 4H₂ → CH₄ (methane under special conditions); N₂ + 3H₂ → 2NH₃ (requires catalyst).
• Reaction with metals: Non-metals react with metals to form ionic or partly ionic compounds (salts). Example: 2Na + Cl₂ → 2NaCl; Cu + Cl₂ → CuCl₂ (on heating). This is often a redox process where the metal is oxidised and the non-metal reduced.
• Reaction with water: Most non-metals do not react directly with water (e.g., C, S, P) under normal conditions. Some non-metals or their compounds react: Cl₂ + H₂O ⇌ HCl + HClO (chlorine reacts with water to form hydrochloric and hypochlorous acids).
• Oxidising and reducing behaviour: Many non-metals (especially oxygen and halogens) are good oxidising agents because they accept electrons. Hydrogen and carbon can act as reducing agents in some reactions.
• Displacement reactions among non-metals (halogens): A more reactive halogen can displace a less reactive halide from its solution: Cl₂ + 2KBr → 2KCl + Br₂. Reactivity order: F > Cl > Br > I.

Important points for Class 8

• Non-metal oxides are generally acidic (they form acids with water).
• Non-metals usually do not displace hydrogen from acids (unlike metals).
• Non-metals form covalent compounds with each other and ionic compounds with metals.
• Many non-metals act as oxidising agents (e.g., oxygen, halogens).

Practical implications
The acidic nature of some non-metal oxides leads to environmental effects such as acid rain (from SO₂, NO₂). Non-metals like chlorine are used for disinfection and bleaching because of their oxidising power. Carbon dioxide from combustion affects air composition and is involved in the greenhouse effect.

Definition: Oxides are binary compounds of oxygen with other elements (general formula: XxOy). They form when elements react with oxygen.

Formation: Metals and non-metals combine with oxygen during burning, heating or chemical reactions. Example reactions: 2Mg + O2 → 2MgO (magnesium oxide), C + O2 → CO2 (carbon dioxide).

Classification of oxides

• Basic oxides: Formed mainly by metals (especially alkali and alkaline earth metals). They react with acids to give salts and water and often dissolve in water to give basic solutions. Example: Na2O + H2O → 2NaOH; CuO + 2HCl → CuCl2 + H2O.
• Acidic oxides: Formed mainly by non-metals. They react with bases to form salts and often form acids on reaction with water. Example: CO2 + H2O → H2CO3; SO2 + 2NaOH → Na2SO3 + H2O.
• Amphoteric oxides: Oxides that react with both acids and bases to form salts and water. Typically formed by some metals (e.g., Al, Zn, Pb). Example: ZnO + 2HCl → ZnCl2 + H2O and ZnO + 2NaOH + H2O → Na2[Zn(OH)4] (or Zn(OH)4^2− solution).
• Neutral oxides: Oxides that do not show acidic or basic behaviour. They do not react appreciably with acids or bases. Examples: CO, N2O, NO.

Properties & tests

• Basic oxides turn red litmus paper blue; acidic oxides turn blue litmus red.
• Reaction with water: Some oxides dissolve in water to give acids or bases (e.g., SO2 → sulfurous acid; Na2O → NaOH), but many oxides (e.g., many metal oxides) are insoluble.
• Reaction with acids/bases: Basic oxides + acid → salt + water. Acidic oxides + base → salt + water. Amphoteric oxides react with both.
• Chemical behaviour often correlates with element's metallic character: more metallic → oxide more basic; more nonmetallic → oxide more acidic.
• Oxidation states: Oxides show the oxidation states of the element combined with oxygen (oxygen is usually −2). Example: Fe2O3 (Fe is +3), CO2 (C is +4).

Uses & importance (real life)

• CaO (quicklime) and Ca(OH)2 (slaked lime) are used in construction, water treatment and manufacture of cement.
• CO2 is used in fire extinguishers, carbonated drinks and as a greenhouse gas in atmosphere.
• Fe2O3 (iron(III) oxide) is rust — important in corrosion engineering; Fe3O4 used as iron ore.
• SiO2 (silicon dioxide) is sand — used in glass, ceramics and electronics.

Summary: Oxides are grouped by chemical behaviour into basic, acidic, amphoteric and neutral. Their reactions with water, acids and bases, along with litmus tests, help classify them. Metallic character of the element largely determines oxide type.

What is the reactivity series?

The reactivity series (or activity series) of metals is a list of metals arranged in order of their chemical reactivity from the most reactive to the least reactive. It helps predict how a metal will react with water, acids, oxygen and salts of other metals.

How the series is constructed

• Metals are arranged by observing their reactions with cold water, steam, dilute acids and by single‑displacement (replacement) reactions with solutions of other metal salts.
• A metal that displaces another metal from its salt solution is placed above that metal in the series.
• Hydrogen is included as a reference point: metals above hydrogen react with dilute acids to give hydrogen gas; those below do not.

Typical CBSE order (most to least reactive)

K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > (H) > Cu > Hg > Ag > Au

Important points and significance

• Metals higher in the series are more reactive: they react more vigorously with water and acids and displace metals lower in the series from solutions of their salts.
• Metals above hydrogen react with dilute acids to produce hydrogen gas.
• Highly reactive metals (e.g., K, Na, Ca) are not found free in nature; they occur as compounds and are often extracted by electrolysis. Less reactive metals (e.g., Cu, Ag, Au) may occur in native (elemental) form.
• Reactivity determines methods of extraction: very reactive metals require electrolysis (e.g., Na, K, Al); moderately reactive metals are reduced by carbon or carbon monoxide in furnaces (e.g., Fe); least reactive metals can often be found as native elements or obtained by simple chemical methods.
• Reactivity series explains practical processes such as corrosion protection (galvanization), sacrificial anodes, and metal displacement reactions used in metallurgy and recycling.

How to test experimentally

• Observe reaction with cold water (very reactive metals): e.g., K, Na react vigorously with cold water producing hydroxides and H2.
• React metals with steam (less reactive metals like Mg react with steam to give oxide and H2).
• React metals with dilute acids (metals above H produce H2).
• Perform displacement reactions: add a metal to a solution of another metal salt; if the solid metal is higher in the series it will replace the metal ion from solution.

Practical consequences

• Galvanization: coating iron with zinc (zinc is above iron and protects it by corroding first - sacrificial protection).
• Sacrificial anodes: attaching a piece of a more reactive metal (e.g., magnesium or zinc) to protect ships, pipelines or underground tanks.
• Choice of extraction method and storage/handling: very reactive metals are stored under oil or in inert conditions to prevent reaction with air or water.

What is extraction of metals? Extraction of metals from ores is the process of obtaining pure metals from naturally occurring mineral rocks (ores) that contain the metal in combined form. The method used depends on the nature of the ore and the reactivity of the metal.

Main stages in extraction

1. Concentration (or dressing) of ore: Removal of impurities (gangue) from the ore by physical methods such as washing, magnetic separation, levigation or froth flotation. This increases the percentage of the metal compound in the sample.
2. Conversion to oxide: Many ores are sulphides or carbonates. They are converted to oxides by roasting (heating in air; sulphides → oxides + SO2) or calcination (heating in absence of air; carbonates → oxides + CO2).
3. Reduction of oxide to metal: The metal oxide is reduced to metal. Methods include:
   • Reduction by carbon or carbon monoxide (for less reactive metals): e.g. ZnO + C → Zn + CO
   • Reduction by more reactive metals (displacement): e.g. Zn + CuSO4 → ZnSO4 + Cu
   • Electrolytic reduction (electrolysis) for very reactive metals like aluminium.
4. Refining / Purification: The extracted metal is often impure and is purified by methods such as electrolytic refining, distillation, zone refining, or chemical methods.

Special methods for some metals

• Blast furnace for iron: Iron is extracted from haematite or magnetite using coke in a blast furnace. Several reactions (combustion of coke, formation of CO, reduction of iron oxide) occur simultaneously. Slag (impurities) is removed using limestone.
• Electrolytic method for aluminium: Bauxite is purified to give alumina (Al2O3). Alumina is dissolved in molten cryolite and reduced by electrolysis to give aluminium metal.
• Hydrometallurgy and leaching: Low-grade ores may be treated with chemical solutions (e.g., cyanide leaching for gold) to dissolve the metal and then recover it from solution.
• Amalgamation and liquation: Special processes used for metals like gold (amalgamation with mercury) and for low melting point metals (liquation).

Why method depends on reactivity: A metal higher in the reactivity series (like aluminium) cannot be reduced by carbon and needs electrolysis. Less reactive metals (like iron, copper, zinc) can be reduced by carbon or carbon monoxide.

Environmental and practical notes: Roasting and reduction produce gases (SO2, CO2) and waste (slag). Modern metallurgy includes measures to control pollution and recycle slag and by-products.

Definition: Corrosion is the gradual chemical or electrochemical destruction of metals caused by reactions with substances in their environment. The most common example is the rusting of iron.

How corrosion (rusting of iron) occurs:

• Rusting is an electrochemical process that needs metal, oxygen and moisture (water). Impurities or salts in water accelerate the process by making the water a better electrolyte.
• Stepwise process:
  • Anodic reaction (iron loses electrons): Fe → Fe2+ + 2e−
  • Cathodic reaction (oxygen gains electrons) in neutral/alkaline water: O2 + 2H2O + 4e− → 4OH−
  • Fe2+ reacts with OH− to form iron(II) hydroxide: Fe2+ + 2OH− → Fe(OH)2
  • Fe(OH)2 is further oxidized by oxygen to form hydrated iron(III) oxide (rust): 4Fe(OH)2 + O2 → 2Fe2O3·nH2O + 2H2O

Factors affecting corrosion: presence of water and oxygen, salt or acidic conditions (lower pH), temperature, stray electrical currents, presence of dissimilar metals in contact (galvanic corrosion), and mechanical stress or cracks.

Types of corrosion (brief): uniform corrosion (surface-wide), galvanic corrosion (between two different metals), pitting (localized holes), crevice corrosion, and intergranular corrosion.

Prevention methods:

• Coating: painting, enamel, plastic coatings or varnish to isolate metal from air and water.
• Oiling and greasing: forms a water-repellent film (common for tools and machine parts).
• Galvanization: coating iron/steel with a layer of zinc. Zinc is more reactive and acts as a sacrificial metal.
• Electroplating/tinning: coating with a less reactive metal (e.g., tin on iron) to protect surface.
• Alloying and passivation: making stainless steel (with Cr) which forms a protective chromium oxide layer; aluminium forms a protective aluminium oxide layer naturally.
• Cathodic protection: connecting the structure to a more easily corroded metal (sacrificial anode) or using impressed current to make the structure the cathode.
• Corrosion inhibitors: chemicals added to the environment (e.g., in cooling water) that slow corrosion reactions.
• Design measures: avoid crevices, ensure drainage, insulate dissimilar metals, and control environment (dehumidification, desalination).

Key idea: Prevention either stops the electrochemical reactions (isolation, inhibitors, drying) or provides a sacrificial path (zinc/ magnesium anodes, cathodic protection).

Definition: An alloy is a homogeneous mixture (usually solid) of two or more elements, where at least one element is a metal. Alloys are made to improve the properties of the base metal — for example, to increase strength, hardness, corrosion resistance or to lower the melting point.

How alloys are made:

• Most common method: the constituent metals are melted together and mixed; the mixture is then cooled to form a solid alloy.
• Some alloys form by diffusion in the solid state (heat treatment), or by rapid solidification.

Types of alloy structures:

• Substitutional solid solutions — atoms of the solute replace atoms of the solvent metal (e.g., brass: Zn atoms replace some Cu atoms).
• Interstitial solid solutions — small atoms fit into the spaces (interstices) between larger metal atoms (e.g., steel: small C atoms in Fe lattice).
• Intermetallic compounds — specific stoichiometric compounds with distinct crystal structures (e.g., some Ni-Al compounds).

Why alloys are useful (property improvements):

• Greater strength and hardness (e.g., steel vs pure iron).
• Improved corrosion resistance (e.g., stainless steel vs iron).
• Lower or controlled melting point (e.g., solders).
• Better wear resistance, electrical or magnetic properties (depends on composition).

Important notes for students: Alloys usually do not show the properties of their constituent pure metals in a simple additive way; sometimes small amounts of an element (like carbon in iron) dramatically change properties. The composition is commonly described by percent by mass (mass percent), mole fraction, or volume fraction.

Examples of common alloys and typical composition/uses:

• Steel: iron + carbon (carbon ≈ 0.02%–2.1% by mass). Used for construction, tools, machines.
• Stainless steel: Fe + C + chromium (Cr ≥ ~12%) ± nickel. Used for cutlery, kitchenware, surgical instruments.
• Brass: copper + zinc (variable Zn 5%–45%). Used for musical instruments, fittings.
• Bronze: copper + tin (typically 12% Sn). Used for statues, bearings, coins.
• Solder (traditional): tin + lead (e.g., 60:40 Sn:Pb). Used to join electrical components (lead-free solders are now common).
• Nickel-chromium (nichrome): Ni + Cr. Used in heating elements.
• Duralumin: aluminium + copper (with Mg, Mn). Used in aircraft structures.

Safety & environmental note: Some alloys (e.g., lead-containing solders, mercury amalgams) are toxic, so safer alternatives and proper disposal are important.

Study tips: Remember the difference between an alloy and a mixture: alloys are homogeneous at macroscopic scale and often have different crystal structures than constituent metals. Learn a few common alloy compositions and their uses (steel, stainless steel, brass, bronze, solder).

Overview
This topic summarizes the common, important chemical reactions involving metals and non-metals that appear in Class 8 (Chapter: Materials: Metals and Non‑metals). Reactions are grouped by type (with oxygen, water, acids, salt solutions, etc.). For each type you will see the general pattern, typical observations, and short explanations.

1. Reactions with oxygen (combustion/oxidation)
Metals and non‑metals burn or react with oxygen to form oxides. Observations: bright flame for some metals (e.g., Mg), heat and light, formation of solid or gaseous oxides.

• General (metal): Metal + O₂ → Metal oxide
• Example: 2Mg + O₂ → 2MgO (white ash; exothermic)
• General (non‑metal): Non‑metal + O₂ → Non‑metal oxide
• Example: C + O₂ → CO₂ (carbon dioxide)

2. Reactions with water
Only reactive metals react with cold water to give metal hydroxides and hydrogen. Less reactive metals react with steam to give metal oxides and hydrogen.

• Cold water (very reactive metals): 2Na + 2H₂O → 2NaOH + H₂
• Steam (less reactive metals): Zn + H₂O (steam) → ZnO + H₂
• Iron with steam: 3Fe + 4H₂O (steam) → Fe₃O₄ + 4H₂

3. Reactions with acids
Most metals (above hydrogen in the reactivity series) react with dilute acids to form a salt and hydrogen gas. Observation: effervescence (bubbles), heat.

• General: Metal + Acid → Salt + H₂
• Example: Zn + 2HCl → ZnCl₂ + H₂

4. Displacement reactions (metal + salt solution)
A more reactive metal displaces a less reactive metal from its salt solution. This is the basis of the activity series.

• Example: Zn + CuSO₄ → ZnSO₄ + Cu (zinc displaces copper)

5. Metal oxides and non‑metal oxides with water
Metal oxides are generally basic (form bases on reacting with water). Non‑metal oxides are generally acidic (form acids with water).

• Basic oxide: Na₂O + H₂O → 2NaOH
• Amphoteric oxide (example): ZnO + H₂O → Zn(OH)₂ (and reacts with acids/bases)
• Acidic oxide: CO₂ + H₂O → H₂CO₃ (carbonic acid)
• SO₂ + H₂O → H₂SO₃ (sulphurous acid)

6. Corrosion and tarnishing
Corrosion is the chemical attack of metals (e.g., iron rusting). Tarnishing is surface reaction (e.g., silver with sulfur).

• Iron rusting (simplified): 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ (which dehydrates to hydrated iron(III) oxide, rust)
• Silver tarnish (simplified): 2Ag + S → Ag₂S (black tarnish)

7. Reduction of metal oxides (extraction of metals)
Some metal oxides can be reduced to the metal by carbon (coke) or by heating. This is used in metallurgy.

• Example: CuO + C → Cu + CO₂ (reduction by carbon)
• Iron oxide reduction in blast furnace (simplified): Fe₂O₃ + 3C → 2Fe + 3CO (industrial conditions)

Observations and how to identify reaction types
Look for gas evolution (hydrogen, CO₂, SO₂), color change, temperature change, flame, precipitate or metal deposition (displacement). Knowledge of activity series helps predict displacement and reactions with water/acids.

Tips to remember
Memorize a short activity series (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au) — metals above H react with acids to give H₂. Metal oxides → bases; non‑metal oxides → acids.

Overview

Metals and non-metals are two broad categories of elements with distinct physical and chemical properties. Their properties determine where they are used in daily life, industry and technology.

Physical and chemical properties (brief)

• Metals: Usually solids (except mercury), lustrous, good conductors of heat and electricity, malleable, ductile, sonorous, generally high melting points and densities. Chemically they tend to lose electrons to form cations and form basic oxides.
• Non-metals: Can be solids, liquids or gases; often dull, poor conductors of heat and electricity, brittle if solid, lower melting points and densities (many). Chemically they tend to gain/share electrons and form acidic or neutral oxides.

Why properties determine uses

The excellent electrical and thermal conductivity, mechanical strength (malleability, ductility) and ability to form alloys make metals ideal for structural, electrical and mechanical applications. The chemical behavior (such as reactivity with oxygen or acids) also affects their selection and protective treatments (e.g., galvanizing iron).

Uses of metals

• Construction and structural: Iron and steel (alloy of iron and carbon) are used in buildings, bridges, ships and railways because they are strong and durable.
• Electrical wiring and electronics: Copper and aluminium are used for electrical wires, motors and transformers due to high electrical conductivity (copper better; aluminium lighter).
• Transport: Aluminium and steel are used in aircraft, automobiles and rail vehicles for strength and weight advantages.
• Household utensils and cookware: Stainless steel, aluminium and copper are used for vessels, pans and cutlery because they conduct heat and are malleable.
• Coins, jewellery and decoration: Gold, silver, platinum and some copper alloys are used for coins and ornaments because of lustre, ductility and resistance to corrosion (gold).
• Batteries and electrochemistry: Metals like zinc, lead and lithium are used as electrodes in batteries.
• Protective coatings and sacrificial anodes: Zinc is used for galvanizing iron to prevent rust; magnesium or zinc blocks protect ship hulls and pipelines.
• Catalysts and high-temperature uses: Platinum and nickel are used as catalysts; tungsten and chromium in high-temperature or wear-resistant applications.

Uses of non-metals

• Respiration and combustion: Oxygen is essential for respiration and for combustion in steel and chemical industries.
• Fertilizers and industrial gases: Nitrogen is used to make ammonia and fertilizers; it is also used as an inert gas in food packaging and industry.
• Disinfectants and bleaching: Chlorine is used for water purification, disinfectants and production of PVC.
• Acids and chemicals: Sulfur is used to make sulfuric acid (important industrial chemical). Phosphorus is used in fertilizers and matches.
• Carbon compounds and fuels: Carbon (in forms like graphite and diamond) is used in pencils, electrodes (graphite) and cutting tools (diamond); hydrocarbons (organic non-metals) are fuels and feedstocks for plastics.
• Insulating materials: Non-metals (rubber, plastics, glass) are used as electrical insulators and protective coatings.
• Medical and household uses: Iodine for antiseptics, helium as a safe inert gas for balloons/cryogenics, oxygen cylinders in medicine.

Important practical notes

• Alloys (mixtures of metals) combine properties: e.g., bronze (Cu + Sn) for statues; brass (Cu + Zn) for musical instruments; stainless steel (Fe + Cr + Ni) resists rusting.
• Corrosion of metals (e.g., rusting of iron) is a major practical problem; protection includes painting, galvanizing, alloying and sacrificial protection.
• Certain non-metals (e.g., chlorine, oxygen at high concentration) can be hazardous and must be handled with care.

Summary table (short)

• Metals: construction, wiring, machines, coins, cookware, batteries, protective coatings.
• Non-metals: respiration, fertilizers, disinfectants, acids, insulators, medicines, fuels and plastics.