Metals And Non Metals Complete - Class 10 Boards Chemistry

Metals and Non-Metals: Properties, Reactions, and Extraction

Physical Properties of Metals and Non-Metals

Metals and non-metals differ distinctly in their physical properties, which are frequently tested in exams. Most properties have exceptions that must be memorized for board questions.

• Lustre: Metals are lustrous (shiny surface, e.g., gold, silver, iron). Non-metals are generally non-lustrous except iodine and graphite.
• Hardness: Metals are generally hard; exceptions are alkali metals (lithium, sodium, potassium, gallium, cesium) which are soft and can be cut with a knife. Non-metals are generally soft; diamond (a form of carbon) is the hardest known natural substance.…

Metals and Non-Metals: Properties, Reactions, and Extraction

Physical Properties of Metals and Non-Metals

Metals and non-metals differ distinctly in their physical properties, which are frequently tested in exams. Most properties have exceptions that must be memorized for board questions.

• Lustre: Metals are lustrous (shiny surface, e.g., gold, silver, iron). Non-metals are generally non-lustrous except iodine and graphite.
• Hardness: Metals are generally hard; exceptions are alkali metals (lithium, sodium, potassium, gallium, cesium) which are soft and can be cut with a knife. Non-metals are generally soft; diamond (a form of carbon) is the hardest known natural substance.
• Malleability: Metals can be beaten into thin sheets (most malleable: gold, silver). Non-metals are non-malleable.
• Ductility: Metals can be drawn into thin wires (e.g., copper, gold). Non-metals are non-ductile.
• Conductivity: Metals are good conductors of electricity and heat (best: silver, copper, gold). Non-metals are poor conductors; graphite conducts electricity.
• Sonority: Metals produce sound when struck (sonorous). Non-metals are non-sonorous.
• Physical State at Room Temperature: All metals are solid except mercury (liquid). Non-metals exist as solids (carbon, sulfur), liquid (bromine), and gases (nitrogen, oxygen, chlorine, fluorine, hydrogen).
• Melting and Boiling Points: Metals generally have high melting/boiling points; exceptions are gallium and cesium (melt on palm). Non-metals generally have low melting/boiling points; diamond is an exception (high).
• Flame Test Colors:
  • Lithium: Crimson (dark red)
  • Sodium: Yellow
  • Potassium: Lilac (light purple)
  • Calcium: Orange-red
  • Copper: Green

Exceptions in Properties

Several exceptions to general trends are frequently asked in exams.

• Soft Metals: Alkali metals (Li, Na, K, Ga, Cs) are soft and can be cut with a knife.
• Low Melting Point Metals: Gallium and cesium melt on the palm.
• Poor Conductors: Lead and mercury are poor conductors of heat; lead is also a poor conductor of electricity.
• Lustrous Non-Metals: Iodine and graphite are lustrous.
• Hard Non-Metal: Diamond (form of carbon) is the hardest natural substance.
• Conducting Non-Metal: Graphite conducts electricity.
• Physical State Exceptions: Mercury (metal, liquid), bromine (non-metal, liquid).
• Corrosion Resistance: Gold and silver are corrosion-resistant, making them suitable for ornaments.

Chemical Properties: Reactions with Oxygen, Water, and Acids

Metals and non-metals show characteristic chemical reactions, with specific products and exceptions.

• Reaction with Oxygen:
  • Metals form metal oxides, generally basic in nature (turn red litmus blue).
  • Non-metals form acidic oxides (e.g., SO₂, CO₂).
  • Potassium and sodium react vigorously, catch fire, and are stored in kerosene.
  • Magnesium burns with dazzling white light to form MgO.
  • Aluminium forms Al₂O₃ with a protective oxide layer.
  • Copper forms black CuO on heating.
  • Gold and silver do not react with oxygen even at high temperature.
• Amphoteric Oxides: Some metal oxides (Al₂O₃, ZnO) react with both acids and bases to form salt and water.
  • With acid: $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$
  • With base: $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O$
• Reaction with Water:
  • Highly reactive metals (K, Na, Ca) react with cold water to form hydroxides and hydrogen gas; K and Na reactions are highly exothermic and hydrogen catches fire.
  • Ca reacts less violently, does not catch fire.
  • Mg reacts with hot water.
  • Al, Zn, Fe react only with steam to form oxides and hydrogen.
  • Metals below iron (Pb, Cu, Ag, Au) do not react with water.
• Reaction with Acids:
  • Metals above hydrogen in reactivity series displace hydrogen from acids (HCl, H₂SO₄) to form salt and hydrogen gas.
  • Copper (below hydrogen) does not react with HCl or H₂SO₄.
  • Hydrogen gas test: colorless, burns with a pop sound, extinguishes burning candle.
  • More reactive metals produce higher temperature and faster hydrogen evolution.
  • With nitric acid (HNO₃), hydrogen gas is not evolved (except Mg and Mn with very dilute HNO₃) due to its strong oxidizing nature.

Reactivity Series and Displacement Reactions

The reactivity series ranks metals by their reactivity, determining their ability to displace others in reactions.

• Reactivity Series Mnemonic: K Na Ca Mg Al Zn Fe Pb H Cu Hg Ag Au
• Displacement Rule: A more reactive metal displaces a less reactive metal from its compound.
• Examples:
  • Fe + CuSO₄ → FeSO₄ + Cu (blue to green solution, copper deposits on iron nail)
  • Zn + CuSO₄ → ZnSO₄ + Cu (blue to colorless solution)
  • Zn + FeSO₄ → ZnSO₄ + Fe (green to colorless solution)
  • Cu cannot displace Zn from ZnSO₄.
• Thermite Reaction: Fe₂O₃ + 2Al → 2Fe + Al₂O₃ (highly exothermic, molten iron used for welding railway tracks and machine parts).

Electronic Configuration and Formation of Ionic Compounds

Metals and non-metals have distinct valence electron configurations, leading to ionic bond formation.

• Metals: 1, 2, or 3 electrons in valence shell; lose electrons to achieve stable octet.
• Non-Metals: 4, 5, 6, or 7 electrons in valence shell; gain electrons to complete octet.
• Ionic Compound Formation:
  • Metal loses electrons (forms cation, e.g., Na⁺, Mg²⁺).
  • Non-metal gains electrons (forms anion, e.g., Cl⁻, O²⁻).
  • Oppositely charged ions attract via strong electrostatic force (ionic/electrovalent bond).
• Examples:
  • NaCl: Na (2,8,1) loses 1 electron, Cl (2,8,7) gains 1 electron.
  • MgO: Mg (2,8,2) loses 2 electrons, O (2,6) gains 2 electrons.
  • MgCl₂: Mg loses 2 electrons, each Cl gains 1 electron (requires 2 Cl atoms).

Properties of Ionic (Electrovalent) Compounds

Ionic compounds exhibit characteristic physical and chemical properties due to strong ionic bonds.

• Hardness and Brittleness: Hard, brittle solids due to strong electrostatic attraction.
• High Melting/Boiling Points: Require high energy to break ionic bonds.
• Solubility: Generally soluble in water, insoluble in organic solvents (alcohol, kerosene, petrol).
• Electrical Conductivity:
  • Do not conduct electricity in solid state (ions fixed).
  • Conduct in molten state and aqueous solution (ions free to move).
• Ionic compounds exist as a combination of ions, not molecules (e.g., NaCl is a lattice of Na⁺ and Cl⁻ ions).

Extraction and Refining of Metals

Metals are extracted from ores using different methods based on their reactivity.

• Minerals: Naturally occurring compounds in earth’s crust.
• Ores: Minerals with high metal content, profitable for extraction.
• Gangue: Impurities (soil, sand) in ores.
• Enrichment/Concentration: Removal of gangue from ore.
• Extraction Methods:
  • Top of Reactivity Series (K, Na, Ca, Mg, Al): Found only in combined state (chlorides, oxides). Cannot be reduced by carbon; extracted by electrolytic reduction.
  • Middle of Reactivity Series (Zn, Fe, Pb): Found as carbonates, oxides, sulfides. Sulfides/carbonates converted to oxides (sulfides by roasting in air, carbonates by calcination in limited air), then reduced by carbon.
    • Roasting: Heating sulfide ore in excess air (ZnS + O₂ → ZnO + SO₂).
    • Calcination: Heating carbonate ore in limited air (ZnCO₃ → ZnO + CO₂).
    • Reduction: ZnO + C → Zn + CO.
  • Bottom of Reactivity Series (Cu, Hg, Ag, Au): Found in free state and combined state. Sulfide ores roasted to oxides, then heated to obtain metal (no reducing agent needed).
    • HgS (cinnabar) → HgO (roasting) → Hg (heating).
    • Cu₂S → Cu₂O (roasting), then Cu₂O + Cu₂S → Cu + SO₂.
• Refining: Metals obtained are impure; purified by electrolytic refining.
  • Anode: Impure metal rod.
  • Cathode: Pure metal strip.
  • Electrolyte: Metal salt solution (e.g., CuSO₄ + dilute H₂SO₄ for copper).
  • On passing current, pure metal deposits on cathode, insoluble impurities settle as anode mud.

Corrosion and Its Prevention

Corrosion is the deterioration of metals due to environmental attack, with specific examples and prevention methods.

• Corrosion: Metal attacked by moisture, oxygen, acids, etc., forming undesirable compounds.
• Rusting of Iron: Formation of hydrated ferric oxide (reddish-brown, flaky) when iron reacts with oxygen and water.
  • Conditions: Iron, oxygen, and water must be present.
• Tarnishing of Copper: Formation of green basic copper carbonate due to moist CO₂.
• Tarnishing of Silver: Formation of black silver sulfide due to atmospheric sulfides.
• Prevention Methods:
  • Painting, oiling, greasing (forms protective layer).
  • Galvanizing: Coating iron/steel with zinc (zinc is more reactive, protects iron even if layer cracks).
  • Anodizing: Forming protective aluminium oxide layer on metals.

Alloys: Composition and Properties

Alloys are homogeneous mixtures of metals or metals with non-metals, with properties different from their constituents.

• Definition: Homogeneous mixture of metals/metals with non-metals.
• Examples:
  • Steel: Iron + 0.05% carbon (harder than pure iron).
  • Stainless Steel: Iron + nickel + chromium (hard, rust-resistant).
  • Amalgam: Alloy containing mercury.
  • Brass: Copper + zinc.
  • Bronze: Copper + tin.
  • Solder: Tin + lead (low melting point).
• Properties: Alloys have properties different from constituent metals (e.g., brass and bronze are poor conductors despite copper being a good conductor; solder melts easily though tin and lead have high melting points).