CBSE Class 10 Science Metals And Non Metals Notes Set 06

These CBSE Class 10 Science notes for Chapter 3, Metals and Non-metals, cover the physical and chemical properties of metals and non-metals, the reactivity series, ionic bonding and the properties of ionic compounds, the occurrence and extraction of metals, electrolytic refining, corrosion and its prevention, and alloys, with the important equations written out and a quick revision summary at the end.

Important Notes and Formulas

Metals, Non-metals and Metalloids

Everything around us is made up of elements; 118 elements are known. On the basis of their properties they are broadly divided into metals and non-metals. Some elements show the properties of both and are called metalloids (boron, silicon, germanium, arsenic, antimony, tellurium). Most of the known elements are metals. All metals are solids except mercury, which is a liquid at room temperature. Among non-metals, bromine is a liquid, some are solids (carbon, sulphur, phosphorus, iodine) and some are gases (hydrogen, oxygen, nitrogen, chlorine).

In the periodic table the metals lie on the left-hand side and in the centre, and the non-metals lie on the right-hand side. A zig-zag line separates them, and the metalloids lie close to this line. Hydrogen is a non-metal placed on the left-hand side.

Physical Properties of Metals

1. Hardness: Most metals are hard and cannot be cut with a knife. Sodium and potassium are soft and can be cut with a knife.
2. Lustre: Metals in their pure state have a shining surface. Metal surfaces look dull when covered with a layer of oxide, carbonate or hydroxide; rubbing with sand paper removes it.
3. Malleability: Metals can be beaten into thin sheets. Gold and silver are the most malleable. Aluminium is made into foil for packing.
4. Ductility: Metals can be drawn into thin wires. Gold is the most ductile; a wire about 2 km long can be drawn from 1 g of gold.
5. Conduction of heat and electricity: Metals are good conductors because they have free electrons. Silver and copper are the best conductors; lead and mercury are comparatively poor conductors of heat.
6. Melting and boiling points: Metals generally have high melting and boiling points. Gallium and caesium melt on the palm of the hand; sodium and potassium also have low melting points.
7. Sonority: Metals produce a ringing sound when struck with a hard object, so bells and gongs are made of metals.
8. Density: Metals generally have high density, but sodium, potassium and lithium are lighter than water.

Physical Properties of Non-metals

1. Non-metals can be solids, liquids or gases at room temperature (phosphorus is solid, bromine is liquid and oxygen is a gas).
2. They are generally brittle, soft and neither malleable nor ductile; solid non-metals break into pieces when hammered. Diamond, an allotrope of carbon, is the hardest natural substance.
3. They are generally non-lustrous. Iodine and graphite are lustrous.
4. They are bad conductors of heat and electricity. Graphite is a good conductor of electricity.
5. They are non-sonorous and have low melting and boiling points (diamond and graphite are exceptions) and generally low density.

Chemical Properties of Metals

Metal atoms generally have 1, 2 or 3 electrons in their outermost shell. They lose these electrons easily to form positive ions and attain the stable electronic configuration of the nearest noble gas, for example Na → Na⁺ + e⁻ (2, 8, 1 → 2, 8), Mg → Mg²⁺ + 2e⁻ and Al → Al³⁺ + 3e⁻. Metals are therefore called electropositive elements.

1. Reaction with Oxygen

Metal + oxygen → metal oxide. Metal oxides are basic in nature. Metals differ in how readily they react with oxygen.

1. Na and K react at room temperature and catch fire, so they are kept under kerosene: 4Na + O₂ → 2Na₂O.
2. Mg burns on heating with a dazzling white flame: 2Mg + O₂ → 2MgO. Aluminium burns on strong heating: 4Al + 3O₂ → 2Al₂O₃.
3. Iron and copper do not burn but form oxides on heating: 3Fe + 2O₂ → Fe₃O₄ and 2Cu + O₂ → 2CuO (black).
4. Silver and gold do not react with oxygen even at high temperatures.
5. The surfaces of Mg, Al, Zn and Pb get covered with a thin protective layer of oxide at ordinary temperature, which prevents further oxidation.

Basic oxides of Na and K dissolve in water to form alkalis: Na₂O + H₂O → 2NaOH and K₂O + H₂O → 2KOH. Most other metal oxides are insoluble in water.

Amphoteric oxides (Al₂O₃, ZnO) react with both acids and bases to form salt and water:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O    Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

ZnO + 2HCl → ZnCl₂ + H₂O    ZnO + 2NaOH → Na₂ZnO₂ + H₂O

2. Reaction with Water

Metal + water → metal hydroxide (or metal oxide) + hydrogen. Metals react with water at different rates.

3. Reaction with Dilute Acids

Metal + dilute acid → salt + hydrogen. Metals above hydrogen in the reactivity series displace hydrogen from dilute hydrochloric acid and dilute sulphuric acid. For example, Mg + 2HCl → MgCl₂ + H₂, 2Al + 6HCl → 2AlCl₃ + 3H₂, Zn + H₂SO₄ → ZnSO₄ + H₂ and Fe + 2HCl → FeCl₂ + H₂. The rate of bubble formation decreases in the order Mg > Al > Zn > Fe, and copper does not react with these acids. Hydrogen is tested by the ‘pop’ sound with a burning matchstick.

Nitric acid is an exception. It is a strong oxidising agent, so it oxidises the hydrogen produced to water and itself is reduced to oxides of nitrogen; hydrogen gas is not evolved. Only magnesium and manganese give hydrogen with very dilute nitric acid. Gold and platinum are not attacked by any single acid but dissolve in aqua regia, a 3 : 1 mixture of concentrated hydrochloric acid and concentrated nitric acid made fresh.

4. Reaction with Salt Solutions (Displacement Reactions)

A more reactive metal displaces a less reactive metal from the solution of its salt: Metal A + salt solution of B → salt solution of A + metal B. Examples:

Zn + CuSO₄ → ZnSO₄ + Cu (blue colour fades, red-brown copper deposits)    Fe + CuSO₄ → FeSO₄ + Cu (blue changes to pale green)    Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. A less reactive metal does not displace a more reactive metal: Cu + ZnSO₄ → no reaction.

Reactivity Series of Metals

The arrangement of metals in the decreasing order of their reactivities is called the reactivity (activity) series. The more easily a metal loses electrons, the higher its reactivity. The series is: K > Na > Ca > Mg > Al > Zn > Fe > Pb > [H] > Cu > Hg > Ag > Au.

1. A metal higher in the series displaces a metal lower in the series from its salt solution.
2. Metals above hydrogen displace hydrogen from dilute acids; metals below hydrogen (Cu, Hg, Ag, Au) do not.
3. Metals at the top (K, Na, Ca, Mg, Al) are never found free in nature, while gold and platinum are found in the free state.
4. The position in the series decides how a metal is extracted from its ore.

Chemical Properties of Non-metals

1. Non-metals react with oxygen to form acidic or neutral oxides: C + O₂ → CO₂ and S + O₂ → SO₂. The acidic oxides dissolve in water to form acids: CO₂ + H₂O → H₂CO₃ and SO₂ + H₂O → H₂SO₃, which turn blue litmus red. Carbon monoxide (CO) and water are neutral oxides.
2. Non-metals do not react with water. Phosphorus is very reactive in air, so it is kept under water.
3. Non-metals do not displace hydrogen from dilute acids, because they gain electrons and cannot supply electrons to H⁺ ions.
4. Non-metals gain electrons to form negative ions and act as oxidising agents. They combine with hydrogen to form stable covalent hydrides, for example N₂ + 3H₂ → 2NH₃ and H₂ + S → H₂S, and with chlorine to form covalent chlorides such as PCl₃.

Ionic Compounds

Atoms of elements combine to attain the stable electronic configuration of a noble gas (octet, or duplet for helium). A metal loses electrons and a non-metal gains them, so oppositely charged ions are formed. The electrostatic force of attraction between them is the ionic (electrovalent) bond, and the compounds formed are called ionic compounds.

Formation of sodium chloride: Na (2, 8, 1) → Na⁺ (2, 8) + e⁻ and Cl (2, 8, 7) + e⁻ → Cl⁻ (2, 8, 8). Formation of magnesium chloride: Mg → Mg²⁺ + 2e⁻; each chlorine atom gains one electron, so MgCl₂ contains one Mg²⁺ and two Cl⁻ ions.

Properties of Ionic Compounds

1. Physical nature: They are solid and somewhat hard because of the strong force of attraction between the ions; they are brittle.
2. Melting and boiling points: They are high, because a lot of energy is needed to break the strong inter-ionic attraction.
3. Solubility: They are generally soluble in water and insoluble in solvents such as kerosene and petrol.
4. Conduction of electricity: They do not conduct in the solid state because the ions are fixed in the rigid structure. They conduct in the molten state and in aqueous solution, because the ions become free to move.

Occurrence of Metals

The earth’s crust is the main source of metals; sea water contains some soluble salts such as sodium chloride. Aluminium is the most abundant metal in the earth’s crust, followed by iron and calcium. Elements or compounds that occur naturally in the earth’s crust are called minerals. A mineral from which a metal can be extracted conveniently and profitably is called an ore. All ores are minerals, but all minerals are not ores. The earthy impurities (sand, soil, rocky material) in an ore are called gangue.

Very unreactive metals such as gold and platinum are found in the free (native) state; silver and copper occur both free and combined. Reactive metals always occur as compounds.

Extraction of Metals

The process of extracting a metal from its ore and refining it for use is called metallurgy. Its three steps are: (1) concentration (enrichment) of the ore by removing the gangue, (2) conversion of the concentrated ore into the metal, and (3) refining of the impure metal.

Conversion of the Ore into the Metal

1. Metals of low reactivity (Hg, Cu): the oxides are reduced to the metal by heating alone. Cinnabar is roasted to HgO: 2HgS + 3O₂ → 2HgO + 2SO₂, and then HgO is heated: 2HgO → 2Hg + O₂. Copper from Cu₂S: 2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂, then 2Cu₂O + Cu₂S → 6Cu + SO₂.
2. Metals of medium reactivity (Zn, Fe, Pb): these occur as sulphides or carbonates, which are first converted to oxides because it is easier to reduce an oxide. Roasting is heating a sulphide ore strongly in excess air: 2ZnS + 3O₂ → 2ZnO + 2SO₂. Calcination is heating a carbonate ore strongly in a limited supply of air: ZnCO₃ → ZnO + CO₂. The oxide is then reduced with carbon (coke): ZnO + C → Zn + CO.
3. Displacement (thermite) method: a more reactive metal such as aluminium can reduce oxides of metals: 3MnO₂ + 4Al → 3Mn + 2Al₂O₃ + heat and Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat. The reaction with iron(III) oxide is the thermite reaction; the heat produced melts the iron, which is used to weld railway tracks and cracked machine parts.
4. Metals of high reactivity (K, Na, Ca, Mg, Al): their oxides cannot be reduced by carbon because these metals have more affinity for oxygen than carbon. They are obtained by electrolytic reduction of the molten chloride or oxide. For sodium: at the cathode Na⁺ + e⁻ → Na and at the anode 2Cl⁻ → Cl₂ + 2e⁻. Aluminium is obtained by electrolysis of molten alumina (Al³⁺ + 3e⁻ → Al at the cathode).

Electrolytic Refining

The metals obtained by reduction are not pure. Copper, zinc, tin, nickel, silver and gold are refined electrolytically. The impure metal is made the anode, a thin strip of pure metal is the cathode and a solution of a salt of the metal is the electrolyte. On passing current, the metal from the anode dissolves into the electrolyte and an equal amount of pure metal is deposited on the cathode. Soluble impurities go into the solution and insoluble impurities settle below the anode as anode mud.

For copper: at anode Cu → Cu²⁺ + 2e⁻; at cathode Cu²⁺ + 2e⁻ → Cu. The electrolyte is acidified copper sulphate solution.

Corrosion

Corrosion is the slow eating away of a metal by the action of air, moisture or chemicals on its surface, forming oxides, carbonates or sulphides. Rusting is the corrosion of iron: iron slowly changes to a brown flaky substance called rust, hydrated iron(III) oxide, Fe₂O₃·xH₂O. Rusting needs both air (oxygen) and water (moisture).

Activity to show the conditions for rusting: three test tubes with clean iron nails: (A) ordinary water, (B) boiled distilled water with a layer of oil, (C) anhydrous calcium chloride. Only the nail in A rusts, showing that both air and water are necessary.

1. Iron: rusting causes a huge loss every year because it damages bridges, railings, ships and car bodies.
2. Aluminium: a thin, tough layer of aluminium oxide forms on the surface and protects the metal below from further corrosion.
3. Copper: in moist air copper slowly forms a green coating of basic copper carbonate, CuCO₃·Cu(OH)₂.
4. Silver: it reacts with hydrogen sulphide (sulphur compounds) in air to form a black coating of silver sulphide, Ag₂S, which is called tarnishing.

Prevention of Corrosion

1. Painting, oiling or greasing, or varnishing the surface, so that air and moisture cannot reach the metal.
2. Galvanisation: coating iron with a thin layer of zinc. Zinc is more reactive than iron, so the iron is protected even if the zinc layer is scratched.
3. Electroplating iron with tin, nickel or chromium.
4. Making alloys: iron mixed with nickel and chromium gives stainless steel, which does not rust.
5. Using anti-rust solutions (for example alkaline phosphate solutions).

Alloys

An alloy is a homogeneous mixture of two or more metals, or of a metal and a non-metal. It is made by melting the main metal and dissolving the other elements in it in definite proportions, then cooling. The properties of an alloy are different from those of its constituents: alloys are usually harder, stronger and more resistant to corrosion, but have a lower electrical conductivity and a lower melting point. Pure gold (24 carat) is too soft for jewellery, so it is alloyed with silver or copper. An alloy in which one metal is mercury is called an amalgam.

Quick Revision Summary

1. Metals lose electrons and form positive ions; non-metals gain electrons and form negative ions.
2. Metals form basic oxides; non-metals form acidic or neutral oxides; Al₂O₃ and ZnO are amphoteric.
3. The reactivity series is K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au. A more reactive metal displaces a less reactive metal from its salt solution.
4. Metals above hydrogen displace hydrogen from dilute acids; nitric acid does not give hydrogen (except with Mg and Mn).
5. Ionic compounds have high melting and boiling points, dissolve in water and conduct electricity only when molten or dissolved.
6. Highly reactive metals are extracted by electrolysis of their molten compounds, medium reactivity metals by reduction of their oxides with carbon, and low reactivity metals by heating their oxides.
7. Sulphide ores are roasted and carbonate ores are calcined to convert them to oxides.
8. Impure metals are refined by electrolysis: impure metal at the anode, pure metal at the cathode.
9. Rusting needs air and moisture; it is prevented by painting, galvanisation, electroplating and alloying.