Metals and Non-metals
CBSE Class 10 Science — Chapter 3 Comprehensive Notes
1. Introduction
Metals and non-metals are important parts of our daily lives. Metals such as iron and copper have different properties and uses, while non-metals such as oxygen, carbon and sulphur also have different properties and uses.
2. Physical Properties
Physical properties are the characteristics of a substance that can be observed without changing its chemical nature.
(a) Metals
Metals show the following physical properties:
Lustre: Metals generally have a shiny and reflective surface called metallic lustre.
Hardness: Iron and copper are hard at room temperature, whereas sodium is so soft that it can be cut with a knife.
Malleability: The ability of a metal to be beaten into thin sheets is called malleability. Gold and silver are among the most malleable metals.
Ductility: The property of metals by which they can be drawn into thin wires is called ductility. Gold is the most ductile metal; about 2 km of wire can be drawn from 1 g of gold.
Conduction of heat and electricity: Metals are generally good conductors of heat and electricity. The best conductors of heat are silver and copper. Lead and mercury are comparatively poor conductors of heat. Electric wires are coated with polyvinyl chloride (PVC) or a rubber-like material.
Sonority: The property of producing a ringing sound when struck against a hard surface is called sonority. Metals that produce a sound on striking a hard surface are said to be sonorous.
Melting and boiling points: Metals generally have high melting points. However, some metals, such as gallium and caesium, have very low melting points.
Non-metals generally have properties opposite to those of metals, although there are several important exceptions.
Lack of lustre: Most non-metals are dull and do not have a shiny surface. However, iodine is a lustrous non-metal.
Poor conduction: Non-metals are generally poor conductors of heat and electricity. Graphite, an allotrope of carbon, is an important exception and conducts electricity.
Brittleness: Many solid non-metals are brittle. They are neither malleable nor ductile. Sulphur and phosphorus are examples.
Variety of states: Non-metals occur in all three states at room temperature. Oxygen is a gas, bromine is a liquid and sulphur is a solid.
Physical properties are useful for classifying elements, but they cannot be used alone because many exceptions exist. Chemical properties provide a clearer basis for classification.
Summary: Exceptions in Physical Properties
3. Chemical Properties of Metals
Metals can be more clearly distinguished from non-metals on the basis of their chemical properties.
(a) Reaction of Metals with Air
Metals combine with oxygen in the air to form metal oxides.
General Reaction: Metal + O₂ → Metal oxide
Examples:
• 2Cu + O₂ → 2CuO (black copper(II) oxide)
• 4Al + 3O₂ → 2Al₂O₃
The reactivity of metals towards oxygen differs from metal to metal:
Sodium and potassium react so vigorously with air that they may catch fire. Therefore, they are stored under kerosene.
Magnesium, aluminium, zinc and lead develop a thin oxide layer on their surfaces. This layer prevents further oxidation.
Iron filings burn vigorously on heating, although a large piece of iron does not burn easily.
Copper does not burn under normal conditions, but it forms copper oxide when heated.
Silver and gold do not react with oxygen even at high temperatures.
Nature of Metal Oxides
Most metal oxides are basic in nature. Some metal oxides, such as aluminium oxide and zinc oxide, show both acidic and basic behaviour. Such oxides are called amphoteric oxides. They react with both acids and bases to form salt and water.
Examples: Aluminium oxide (Al₂O₃) and Zinc oxide (ZnO).
(b) Reaction of Metals with Water
Some metals react with water to form metal hydroxides or metal oxides and hydrogen gas. However, all metals do not react with water.
• Metal + Water → Metal hydroxide / oxide + Hydrogen gas
• Metal oxide + Water → Metal hydroxide (if soluble)
Reaction with Cold Water
Potassium and sodium react violently with cold water. The reaction is highly exothermic, and the hydrogen produced may catch fire:
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heatCalcium reacts less violently with cold water and starts floating because hydrogen bubbles stick to its surface:
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)Magnesium does not react appreciably with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It may float because of hydrogen bubbles.
Reaction with Steam
Aluminium, iron and zinc do not react appreciably with cold or hot water, but they react with steam to form metal oxides and hydrogen:
• 2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)
• 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
Note: Lead, copper, silver and gold do not react with water under ordinary conditions.
(c) Reaction of Metals with Dilute Acids
General Reaction: Metal + Dilute acid → Salt + H₂
Examples:
• Fe + 2HCl → FeCl₂ + H₂
• Mg + 2HCl → MgCl₂ + H₂
• Zn + 2HCl → ZnCl₂ + H₂
• 2Al + 6HCl → 2AlCl₃ + 3H₂
Note: Copper does not react with dilute hydrochloric acid because copper is below hydrogen in the reactivity series.
Reaction with Nitric Acid (HNO₃): Hydrogen gas is not evolved when a metal reacts with nitric acid because HNO₃ is a strong oxidising agent. It oxidises the hydrogen produced to water and itself gets reduced to nitrogen oxides such as N₂O, NO and NO₂. Magnesium and manganese react with very dilute HNO₃ to evolve hydrogen gas.
(d) Reaction of Metals with Solutions of Other Metal Salts
General Reaction: Metal A + Salt solution of metal B → Salt solution of metal A + Metal B
A more reactive metal can displace a less reactive metal from its salt solution.
• Test tube 1: Fe + CuSO₄ → FeSO₄ + Cu (Iron displaces copper from copper sulphate solution because iron is more reactive than copper).
• Test tube 2: Cu + FeSO₄ → No reaction (Copper cannot displace iron from iron sulphate solution because copper is less reactive than iron).
(e) The Reactivity Series
The reactivity series is a list of metals arranged in the order of their decreasing chemical activity.
Mnemonic: Please Study Carefully, My Aunt Zara Is Learning Hard; Children Must Study Again.
• P: Potassium (K)
• S: Sodium (Na)
• C: Calcium (Ca)
• M: Magnesium (Mg)
• A: Aluminium (Al)
• Z: Zinc (Zn)
• I: Iron (Fe)
• L: Lead (Pb)
• H: Hydrogen (H)
• C: Copper (Cu)
• M: Mercury (Hg)
• S: Silver (Ag)
• A: Gold (Au)
4. How do Metals and Non-metals React?
Reactivity is the tendency of an element to attain a stable electronic configuration. Metals generally lose electrons from their valence shell to form positively charged ions called cations. Non-metals generally gain electrons to form negatively charged ions called anions.
Example 1: Formation of Sodium Chloride (NaCl)
Sodium (Na):
• A sodium atom has one electron in its outermost shell.
• It loses this electron. Its next shell becomes the outermost shell and contains eight electrons, giving a stable configuration.
• A neutral sodium atom has 11 protons and 11 electrons. After losing one electron, it has 11 protons and 10 electrons, becoming Na⁺.
Na → Na⁺ + e⁻
Chlorine (Cl):
• A chlorine atom has seven electrons in its outermost shell and needs one more electron to complete the shell.
• It accepts the electron lost by sodium.
• A neutral chlorine atom has 17 protons and 17 electrons. After gaining one electron, it has 17 protons and 18 electrons, becoming Cl⁻.
Cl + e⁻ → Cl⁻
The oppositely charged sodium and chloride ions attract one another and form sodium chloride. Sodium chloride does not exist as separate molecules. It exists as a three-dimensional arrangement of oppositely charged ions held together by strong electrostatic forces.
Example 2: Formation of Magnesium Chloride (MgCl₂)
Magnesium loses two electrons to form Mg²⁺, while each chlorine atom gains one electron to form Cl⁻. Thus, magnesium chloride contains one Mg²⁺ ion and two Cl⁻ ions.
• Mg → Mg²⁺ + 2e⁻
• 2Cl + 2e⁻ → 2Cl⁻
• Mg²⁺ + 2Cl⁻ → MgCl₂
5. Ionic Compounds
Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.
Properties of Ionic Compounds
Physical nature: Ionic compounds are generally hard crystalline solids because of the strong attraction between oppositely charged ions. They are brittle and may break into fragments when pressure is applied.
Melting and boiling points: Ionic compounds have high melting and boiling points because considerable energy is needed to overcome the strong inter-ionic attractions.
Solubility: Ionic compounds generally dissolve in water but do not dissolve in non-polar solvents such as kerosene or petrol.
Conduction of electricity: Solid ionic compounds do not conduct electricity because their ions are fixed in position. In the molten state or in aqueous solution, the ions become free to move and conduct electricity.
6. Occurrence of Metals
The earth's crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride and magnesium chloride. Some metals are found in nature in the free state, while others are found in the form of their compounds. Gold, silver, platinum and copper are examples of metals found in the free state.
Key Definitions
Minerals: The elements or compounds that occur naturally in the earth's crust are called minerals.
Ores: Minerals that contain a high percentage of a particular metal and from which the metal can be extracted profitably are called ores.
Gangue: The unwanted impurities, such as soil and sand, present in an ore.
Steps in the Extraction of Metals
1. Enrichment of ores: Removal of gangue (unwanted impurities) from the ore based on differences in physical or chemical properties.
2. Extraction of metals: Conversion of concentrated ore into crude metal.
3. Refining of metals: Purification of the extracted metal.
Extraction Methods Based on Reactivity
(i) Extracting Metals Low in the Activity Series (Hg, Cu)
Metals low in the activity series are very unreactive. Their oxides can be reduced to metals by heating alone.
• Mercury from Cinnabar (HgS):
2HgS + 3O₂ —(Heat)→ 2HgO + 2SO₂
2HgO —(Heat)→ 2Hg + O₂
• Copper from Copper Sulphide (Cu₂S):
2Cu₂S + 3O₂ —(Heat)→ 2Cu₂O + 2SO₂
2Cu₂O + Cu₂S —(Heat)→ 6Cu + SO₂
(ii) Extracting Metals in the Middle of the Activity Series (Fe, Zn, Pb, Cu)
Metals in the middle have moderate reactivity. Ores are commonly sulphides or carbonates and are converted into oxides first.
Roasting: Heating sulphide ores strongly in excess air to convert them into oxides.
2ZnS(s) + 3O₂(g) —(Heat)→ 2ZnO(s) + 2SO₂(g)
Calcination: Heating carbonate ores strongly in limited air to convert them into oxides.
ZnCO₃(s) —(Heat)→ ZnO(s) + CO₂(g)
Reduction of Metal Oxides:
• Using Carbon (coke): ZnO(s) + C(s) —(Heat)→ Zn(s) + CO(g)
• Using Reactive Metals (Displacement): 3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + heat
• Thermite Reaction: Reaction of iron(III) oxide with aluminium is highly exothermic. Used for joining railway tracks:
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + heat
(iii) Extracting Metals at the Top of the Activity Series (Na, Mg, Ca, Al)
Highly reactive metals cannot be obtained by heating oxides with carbon because they have a stronger attraction for oxygen than carbon has. They are obtained by electrolytic reduction of molten salts:
• At Cathode (Negative electrode): Na⁺(molten) + e⁻ → Na(s)
• At Anode (Positive electrode): 2Cl⁻(molten) → Cl₂(g) + 2e⁻
Refining of Metals: Electrolytic Refining
The most widely used method for refining crude metal is electrolytic refining (e.g. copper, zinc, tin, silver, gold).
• Anode: Impure metal block
• Cathode: Thin strip of pure metal
• Electrolyte: Acidified solution of the metal salt (e.g. acidified CuSO₄)
• Mechanism: On passing electric current, pure metal dissolves from anode into solution and deposits at cathode. Insoluble impurities settle below the anode as anode mud.
7. Corrosion
Corrosion is the gradual deterioration of metal surfaces when exposed to moist air over prolonged periods.
Conditions&Examples:
1. Iron: Rusting requires BOTH air and water, forming brown flaky rust.
2. Silver: Turns black due to reaction with sulphur in air, forming silver sulphide coating.
3. Copper: Reacts with moist carbon dioxide in air, developing a green coating of basic copper carbonate.
Prevention of Corrosion
Methods include painting, oiling, greasing, galvanising, chrome plating, anodising, or alloying.
• Galvanisation: Protecting iron/steel from rusting by coating with a thin layer of zinc. Zinc protects even if coating is scratched.
• Iron Pillar at Delhi: Built over 1600 years ago near Qutub Minar, displaying remarkable ancient Indian metallurgical rust-resistance.
8. Alloys
An alloy is a homogeneous mixture of two or more metals, or of a metal and a non-metal. It is prepared by melting the primary metal and mixing other elements in suitable proportions.
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