CLASS 10 SCIENCE NOTES: CHAPTER - METALS AND NON - METALS

Metals and Non‑Metals — Class 10 Notes

1. What are Metals and Non‑Metals?

Metals are elements that typically lose electrons to form positive ions (cations). Examples: iron (Fe), copper (Cu), aluminium (Al), silver (Ag), gold (Au).

Non‑metals are elements that usually gain electrons to form negative ions (anions) or share electrons (covalent bonding). Examples: sulphur (S), phosphorus (P), oxygen (O), carbon (C).

2. Physical properties — quick comparison

Metals

Property	Observation / Example	Notes / Exceptions
Appearance (lustre)	Shiny metallic surface (freshly cut)	Iron, Cu, Al
Hardness	Generally hard	Varies — sodium & potassium are very soft
Malleability	Can be beaten into thin sheets	Gold & silver are most malleable
Ductility	Can be drawn into wires	Gold is extremely ductile (1 g → long wire)
Conductivity	Good conductors of heat & electricity	Silver, copper are best conductors; (lead & mercury are poor conductors)
Melting/boiling points	Generally high	Gallium & Cesium melt near room temperature
Sonority	Produce ringing sound when struck	Used in bells

Non‑Metals

Property	Observation / Example	Notes / Exceptions
Occurrence / state	Exist mostly as solids or gases.			Bromine is the only liquid non‑metal (exception)
Appearance	Usually dull (non‑lustrous).	Iodine is (shiny) lustrous (exception)
Hardness	Generally soft	Diamond (allotrope of C) is hardest natural substance (exception)
Malleability / ductility	Brittle — break when hammered	Opposite of metals
Conductivity	Poor conductors of heat & electricity	Graphite conducts electricity (exception)
Melting/boiling points	Usually low	Diamond has very high m.p. & b.p. (exception)
Sonority	Not sonorous (no ringing)	Opposite of metals

Summary table — Metals vs Non‑metals

Property	Metals	Non‑metals
State	Mostly solids at room temp (Mercury liquid)	Solids or gases (Bromine liquid)
Lustre	Metallic, shiny	Usually dull (Iodine exception)
Hardness	Generally hard	Generally soft (Diamond exception)
Malleability	Malleable	Brittle
Ductility	Ductile	Not ductile (Gold exception)
Conductivity	Good conductors	Poor conductors (graphite exception)
Melting/boiling	Usually high	Usually low (diamond exception)
Sonority	Sonorous	Not sonorous

3. Important Applications (and why)

• School bells: Made of metals because metals are sonorous and produce ringing sound.
• Aluminium foil: Used for food packaging — aluminium is malleable, non‑toxic and does not react with food.
• Wires coated with PVC/rubber: PVC & rubber are insulators (non‑metals) that prevent electric shocks by stopping current leakage.
• Copper wires: Copper is ductile and an excellent conductor of electricity, so wires are made from copper.
• Gold & silver for ornaments: Gold and silver are lustrous, malleable, ductile and resistant to corrosion so they stay shiny for long.
• Cooking utensils: Copper and aluminium conduct heat well and have high melting points, so they are suitable for cookware.

Some important questions:

• Which metal can be cut with a knife? 

— Answer: Sodium (Na) or Potassium (K), they are very soft.

• Which is a lustrous non‑metal? 

— Answer: Iodine (I).

• Which metals can melt on the palm due to low m.p.? 

— Answer: Gallium (Ga) and Cesium (Cs).

• Liquid non‑metal? 

— Answer: Bromine (Br).

4. Chemical properties — Metals

Metals show characteristic reactions with oxygen, water and acids. Their position in the reactivity series tells us how they behave.

The Reactivity Series is a list of metals arranged in the order of their decreasing reactivity.

More reactive metals displace less reactive metals from their compounds.

4.1 Reaction with oxygen

Most metals react with oxygen to form metal oxides (basic in nature). 

Examples and equations:

• Potassium: 4K + O₂ → 2K₂O (very vigorous and catches fire if kept in open)
• Sodium: 4Na + O₂ → 2Na₂O (very vigorous and catches fire if kept in open)

NOTE: Sodium and potassium are kept in kerosene oil to prevent accidental fires.

• Magnesium: 2Mg + O₂ → 2MgO (basic oxide) (burns with dazzling white flame)
• Aluminium: 4Al + 3O₂ → 2Al₂O₃ (protective Al₂O₃ layer — amphoteric)
• Iron (filings): 3Fe + 4O₂ → 2Fe₃O₄ (Iron does not burn easily as lump, but filings do) 
• Copper: 2Cu + O₂ → 2CuO (black oxide layer)

• Silver and Gold does not not react even at high temperature.

Some metal oxides (e.g., Al₂O₃, ZnO) are amphoteric as they show both acidic and basic behaviour. Such oxides react with both acids and bases to form salt and water. 

Example reactions:

Reaction with acid(act as a base): Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
Reaction with base(act as an acid): Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

4.2 Reaction with water

Metal + H₂O → Metal oxide + H₂

Metal oxide + H₂O → Metal hydroxide

Reactivity varies with the metal:

• Alkali metals (K, Na): react vigorously with cold water, producing hydroxides and H₂.

2Na + 2H₂O → 2NaOH + H₂ + Heat

Reaction is so violent and exothermic that hydrogen catches fire.

• Ca + 2H­₂O → Ca(OH)­₂ + H­₂                                                                                                                                          • Reacts less violently with cold water.

• • Hydrogen does not catch fire.
• • Calcium floats (due to H­₂ bubbles sticking to surface).

With hot water:

• Mg + 2H­₂O (hot) → Mg(OH)­₂ + H­₂

• • Also floats (H­₂ bubbles stick to surface)
• • No reaction with cold water.

React with steam:

• • 2Al + 3H­₂O (steam) → Al­₂O­₃ + 3H­₂
• • 3Fe + 4H­₂O (steam) → Fe­₃O­₄ + 4H­₂
• • Zn + H­₂O (steam) → ZnO + H­₂

No reaction with cold/hot water.

Do not react with water at all.

• Lead (Pb), Copper (Cu), Silver (Ag), Gold (Au)

4.3 Reaction with dilute acids

Metals react with dilute acids to form corresponding salts and hydrogen gas:

Metal + Dilute acid → Salt + H₂

Examples:

• Mg + 2HCl → MgCl₂ + H₂ 

( Reacts most vigorously, Fastest bubble formation and highest temperature.)

• 2Al + 6HCl → 2AlCl₃ + 3H₂ (Reacts vigorously but less than Mg.)
• Zn + 2HCl → ZnCl₂ + H₂ (Moderate reaction)
• Fe + 2HCl → FeCl₂ + H₂ (slower compared to Mg, Al, Zn)

Exception: Copper does not react with dilute HCl.

• Nitric acid (HNO₃) is a strong oxidising acid: hydrogen gas is not evolved when metals react with nitric acid because HNO₃ oxidises H₂ to water and itself is reduced to nitrogen oxides (NO, NO₂, N₂O). Very dilute HNO₃ can give H₂ with some metals (e.g., Mg).

• With very dilute nitric acid, magnesium (Mg) and manganese (Mn) can evolve hydrogen gas because the acid is too weakly oxidising in such conditions.

4.4 Aquaregia

• Aquaregia is a fresh mixture of concentrated HCl and concentrated HNO₃ in ratio 3:1.
• HCl provides chloride ions and HNO₃ acts as an oxidiser. Nitric acid oxidises metals and gets itself reduced to oxides of Nitrogen such as NO etc.
• It is highly corrosive, fuming liquid capable of dissolving noble metals like gold and platinum which do not dissolve in either acid alone.
• Application: Used in purification and testing of gold and platinum in laboratories and jewellery shops.

4.5 Displacement reactions (metals with metal salt solutions)

Metal A + Salt solution of B → Salt solution of A + Metal B

A more reactive metal displaces a less reactive metal from its salt solution. 

Example:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Here iron displaces copper (Fe more reactive).

5. Chemical properties — Non‑metals

5.1 Reaction with oxygen

Non- metal + Oxygen → Non - metal oxide 

Example:

• C (s) + O­₂ (g) → CO­₂ (g)

• S (s) + O­₂ (g) → SO­₂ (g)

Non‑metals form non‑metallic oxides. These oxides are typically acidic (CO₂, SO₂) or neutral (CO, H₂O).

5.2 Non‑metal oxides + water

Non‑metal oxides react with water to form acids:

• CO₂ + H₂O → H₂CO₃ (carbonic acid)
• SO₃ + H₂O → H₂SO₄ (sulphuric acid)
• 2NO₂ + H₂O → HNO₂ + HNO₃

Generally non-metals do not react with water because they cannot displace hydrogen from water.

6. Electronic configurations & ionic bond formation

• Atoms react to attain stable electronic configuration (octet or duplet). 
• Metals tend to lose electrons (form cations)
• Non‑metals gain electrons (form anions). 
• Electrostatic attraction between cations and anions forms ionic (electrovalent) compounds that involves the transfer of electrons

Example: 

• Formation of Sodium Chloride, NaCl

• Formation of Magnesium Chloride, MgCl₂  

Properties of ionic compounds

• Usually solids, hard (due to strong force of attraction between the oppositely charged ions) and brittle (breaks when high pressure is applied)
• High melting & boiling points (large energy is required to break strong inter- ionic forces).
• Generally soluble in water, insoluble in organic solvents (petrol, kerosene etc.)
• Do not conduct electricity in solid state (ions fixed). Conduct in molten state and in aqueous solution (ions free to move).

7. Reactivity series (activity series)

The reactivity series arranges metals in decreasing order of reactivity. Highly reactive metals (K, Na, Ca, Mg, Al) are at the top; least reactive (Au, Pt) at the bottom. A more reactive metal can displace a less reactive metal from its compounds.

Consequences: Highly reactive metals are never found in native state; they occur as compounds (oxides, carbonates, sulphides). Less reactive metals (Cu, Ag, Au, Pt) can be found free in nature.

• Metallurgy - is the science and technology of extracting metals from their ores, refining them, and preparing them for use.
• Minerals: Naturally occurring elements or compounds present in the earth's crust.
• Ores: Those minerals from which metals can be profitably and conveniently extracted.
• Gangue: The impurities such as sand, soil, and rocky materials that are present along with the ore are called gangue.

Classification of Metals by Reactivity

VERY REACTIVE METALS

• They are never found in free state in nature.

• They always occur in combined state as their oxides, carbonates or sulphides.

MODERATELY REACTIVE METALS

• They are found in the Earth's crust mainly as oxides, sulphides and carbonates.

• Examples:

  • Zinc as ZnS (zinc blende), ZnCO₃ (calamine),

  • Iron as Fe₂O₃ (haematite),

  • Lead  as PbS (galena)

LEAST REACTIVE METALS.

• They are found in free state (native state) as well as in combined state.

• Examples:

  • Copper, Silver, Mercury  found both free and combined (as oxides/sulphides).

  • Gold, Platinum  found in free (native) state.

Note: The ores of many metals are oxides, because oxygen is very reactive and abundant in the earth's crust.

8. Occurrence of metals and metallurgy

Metals occur in the Earth's crust either in combined form (oxides, sulphides, carbonates) or in native/free form. Ores are minerals from which metals can be profitably extracted. Gangue is the unwanted rocky material mixed with the ore.

Steps in metallurgy (overview)

1. Crushing & grinding — reduces ore to powder to make extraction easier and efficient.
2. Concentration (enrichment) — removes gangue (like sand, clay and rocky substances) to increase metal percentage in ore. Done by various physical/ chemical methods depending upon the ore and impurity.
3. Extraction — obtain metal from concentrated ore (method depends on metal reactivity). Achieved by reduction (removal of oxygen or non - metallic element)

• Low reactivity metals (e.g., Hg, Cu, Ag): oxides can be reduced by heating or mild reducing agents.
• Medium reactivity metals (e.g., Fe, Zn, Pb): ores often roasted/calcined to oxides and then reduced by carbon or displacement.
• High reactivity metals (e.g., K, Na, Ca, Mg, Al): obtained by electrolytic reduction of molten salts (electrolysis).

4. Refining — purify the obtained metal (electrolytic refining commonly used).

Roasting vs Calcination

Process	When used	Chemical effect
Roasting	Sulphide ores	Converts MS → MO + SO₂ (e.g., 2ZnS + 3O₂ → 2ZnO + 2SO₂)
Calcination	Carbonate ores	Converts MCO₃ → MO + CO₂ (e.g., ZnCO₃ → ZnO + CO₂)

Thermit reaction (special)

Aluminium reduces iron(III) oxide to produce molten Fe and Al₂O₃ with large heat — used for welding railway tracks:

Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat

Examples of extraction

Mercury from cinnabar (HgS) — heating converts HgS → HgO + SO₂, then HgO → Hg + 1/2 O₂

Copper from Cu₂S — roasting and reduction steps give metallic copper with release of SO₂.

9. Refining of metals

Metals obtained after extraction are impure. Electrolytic refining is widely used (for Cu, Ag, Au, Zn, Sn, Ni):

• Anode: impure metal
• Cathode: thin strip of pure metal
• Electrolyte: solution of metal salt (e.g., CuSO₄ for copper)
• On passing current, pure metal deposits at cathode; soluble impurities go into solution; insoluble impurities form anode mud (may contain Ag, Au).

10. Corrosion and prevention (rusting)

Corrosion is the slow chemical attack of metals by air/moisture/chemicals. Rusting is corrosion of iron forming hydrated iron(III) oxide (Fe₂O₃·xH₂O).

Conditions for rusting

Both oxygen and water (moisture) must be present for rusting. If either is absent, rusting does not occur (demonstrated by experiments using boiled water + oil, or desiccant).

Other corrosion examples

• Silver tarnishes (Ag₂S forms) on exposure to sulphur compounds in air.
• Copper develops a green layer of basic copper carbonate (CuCO₃·Cu(OH)₂).

Methods to prevent rusting

1. Painting, oiling, greasing (barrier methods).
2. Galvanisation (coating iron with zinc). Zinc gives sacrificial protection.
3. Chrome plating (electroplating with chromium).
4. Anodising (form protective oxide layer on Al).

11. Alloying

An alloy is a homogeneous mixture of two or more metals, or a metal and a non‑metal. Alloys are made to improve mechanical properties and corrosion resistance.

• Preparation: Melt primary metal, dissolve other elements in definite proportions, cool.
• Properties: Usually lower electrical conductivity than pure metal, lower m.p., greater strength and resistance to corrosion.

Common alloys

Alloy	Composition	Use / Property
Brass	Cu + Zn	Harder than Cu, used for fittings
Bronze	Cu + Sn	Hard, corrosion resistant
Steel	Fe + C	Harder & stronger than iron
Stainless steel	Fe + Cr + Ni	Resistant to corrosion
Solder	Pb + Sn	Low melting point for joining
Amalgam	One component is Hg (mercury)	Dental amalgam (historic/use varies)

Why 22‑carat gold is used for jewellery: 24‑carat gold is pure but very soft. 22‑carat is an alloy (gold + small Cu/Ag) — harder and more durable for ornaments.

12. Useful practice questions & short answers

• Why are sodium & potassium stored in kerosene? 

— They react vigorously with air & water; kerosene prevents contact with moisture/oxygen.

• Why are electrical wires coated? 

— To insulate and prevent electric shocks; coating is non‑metallic (PVC/rubber).

• Why does calcium float on water? 

— H₂ bubbles adhere to its surface making it buoyant. Reaction: Ca + 2H₂O → Ca(OH)₂ + H₂.

• Why is hydrogen not evolved with nitric acid? 

— HNO₃ oxidises H₂; nitrogen oxides are produced instead. Exception: very dilute HNO₃ with some metals like Mg may give H₂.

Board style Questions:

1. Show formation of CaCl₂ by electron transfer from Ca (Z=20) to Cl (Z=17). Write electronic configuration and formula.
2. Why do ionic compounds conduct electricity in molten state but not in solid state?
3. Describe steps in extraction of copper from its sulphide ore and write chemical equations.