Metals and Non-Metals - Complete CBSE Class 10 Science Notes

Introduction to Metals and Non-Metals

Scientists have discovered over 118 elements, which are broadly classified into metals, non-metals, and metalloids. Understanding the properties and behavior of these elements is fundamental to chemistry and has numerous practical applications in our daily lives.

Metals are electropositive elements that are typically hard, sonorous, malleable, ductile, with high tensile strength and excellent conductivity of heat and electricity. Non-metals exhibit properties opposite to metals, while metalloids show characteristics of both.

Physical Properties: Metals vs Non-Metals

Comparing Physical Characteristics

Understanding the physical differences between metals and non-metals helps identify and classify elements effectively.

Key Examples

Metals:

• Gold (Au) and Silver (Ag): Best conductors of electricity, highly malleable
• Sodium (Na) and Potassium (K): Soft metals that can be cut with a knife
• Iron (Fe): Magnetic properties, high tensile strength
• Aluminium (Al): Light, forms protective oxide layer

Non-Metals:

• Diamond (C): Hardest known substance
• Graphite (C): Conducts electricity, soft and slippery
• Bromine (Br): Only liquid non-metal at room temperature
• Iodine (I): Solid non-metal with metallic lustre

Chemical Properties of Metals

1. Reaction of Metals with Oxygen

Almost all metals combine with oxygen to form metal oxides, which are generally basic in nature.

General Reaction:

Metal + Oxygen → Metal Oxide

Specific Examples:

Amphoteric Oxides: Some metal oxides like Al₂O₃ and ZnO are amphoteric - they react with both acids and bases:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (reaction with acid)Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (reaction with base)

Reactivity Order with Oxygen: Na > Mg > Zn > Fe > Cu

Important Process - Anodising: Anodising forms a thick protective oxide layer on aluminium. The clean aluminium article is made the anode and electrolysed with dilute H₂SO₄. The oxygen evolved at the anode reacts with aluminium to create a thicker protective layer that can be dyed for attractive finishes.

2. Reaction of Metals with Water

Metals react with water to produce metal oxides (or hydroxides) and hydrogen gas. The reactivity varies significantly:

General Reactions:

Metal + Water → Metal Oxide + HydrogenMetal Oxide + Water → Metal Hydroxide

Reactivity Categories:

Reactivity Order with Water: K > Na > Ca > Mg > Zn > Fe > Cu

Safety Note: Sodium and potassium are kept in kerosene to prevent contact with moisture in air, which would cause violent reactions.

3. Reaction of Metals with Acids

Metals react with dilute acids to form metal salts and release hydrogen gas. The metal replaces hydrogen atoms in the acid.

General Reaction:

Metal + Dilute Acid → Metal Salt + Hydrogen

Specific Examples:

Mg + 2HCl → MgCl₂ + H₂↑Zn + 2HCl → ZnCl₂ + H₂↑Fe + 2HCl → FeCl₂ + H₂↑Cu + HCl → No reaction

Reactivity Order with Acids: Na > Mg > Zn > Fe > Cu

Important Notes:

• Hydrogen gas is NOT evolved when metals react with nitric acid (HNO₃), as it's a strong oxidising agent that oxidises H₂ to water
• Exception: Very dilute HNO₃ reacts with Mg and Mn to produce H₂
• Aqua Regia (3 parts HCl + 1 part HNO₃) can dissolve noble metals like gold and platinum

Au + 3Cl → AuCl₃ (in aqua regia)Pt + 4Cl → PtCl₄ (in aqua regia)

4. Reaction with Chlorine

Metals combine with chlorine to form ionic chlorides (electrovalent compounds):

2Na + Cl₂ → 2NaClCa + Cl₂ → CaCl₂Mg + Cl₂ → MgCl₂Zn + Cl₂ → ZnCl₂

During these reactions:

• Metal loses electrons → becomes positively charged (cation)
• Chlorine gains electrons → becomes negatively charged (anion)
• Metal undergoes oxidation
• Chlorine undergoes reduction

5. Reaction with Hydrogen

Only a few reactive metals form ionic hydrides by reacting with hydrogen:

2Na + H₂ → 2NaH (Sodium hydride)2K + H₂ → 2KH (Potassium hydride)Ca + H₂ → CaH₂ (Calcium hydride)

Most metals don't form hydrogen compounds as metals lose electrons while hydrogen typically shares or loses electrons rather than accepting them.

The Reactivity Series of Metals

Understanding Metal Reactivity

The Activity Series (or Reactivity Series) arranges metals in decreasing order of their chemical reactivity. This series is fundamental for predicting chemical reactions.

Complete Reactivity Series

Most Reactive (Top)↓Potassium (K)Sodium (Na)Barium (Ba)Calcium (Ca)Magnesium (Mg)Aluminium (Al)Zinc (Zn)Iron (Fe)Nickel (Ni)Tin (Sn)Lead (Pb)[HYDROGEN (H)] ← Reference pointCopper (Cu)Mercury (Hg)Silver (Ag)Gold (Au)Platinum (Pt)↓Least Reactive (Bottom)

How the Reactivity Series is Determined

The reactivity series is established through various experimental observations:

1. Reaction with Oxygen: More reactive metals combine readily with oxygen
2. Reaction with Water: Reactivity measured by vigor of reaction with cold/hot water or steam
3. Reaction with Acids: Speed and intensity of hydrogen evolution indicates reactivity
4. Displacement Reactions: A more reactive metal displaces a less reactive metal from its salt solution
5. Ionization Energy: Metals with lower ionization energy are more reactive

Displacement Reaction Example:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)(More reactive) (Less reactive displaced)

The blue color of copper sulfate fades as zinc displaces copper, demonstrating zinc's higher reactivity.

Related Helpful Resources

Significance of the Reactivity Series

1. Hydrogen Displacement: Metals above hydrogen in the series can displace hydrogen from dilute acids

Mg + H₂SO₄ → MgSO₄ + H₂ Zn + H₂SO₄ → ZnSO₄ + H₂

2. Metal Displacement: A more electropositive metal replaces a less electropositive metal from solutions

Fe + CuSO₄ → FeSO₄ + Cu (Iron is more electropositive than copper)

3. Extraction Methods: Determines the method required to extract metals from ores
   • Metals at top: Require electrolysis
   • Metals in middle: Reduction with carbon/other reducing agents
   • Metals at bottom: Heating alone or found in native state
4. Corrosion Prediction: More reactive metals corrode faster
5. Electropositive Character: Decreases from potassium to gold

Occurrence and Extraction of Metals

Natural Occurrence

Free State (Native Metals): Metals that remain unaffected by moisture, oxygen, and CO₂ occur in native form:

• Gold (Au)
• Platinum (Pt)
• Silver (Ag)
• Mercury (Hg)

Combined State: Reactive metals occur as compounds in minerals:

• Sodium, Calcium, Potassium
• Aluminium, Lead, Copper, Iron

Minerals and Ores

Mineral: Metal-bearing substances naturally occurring in the earth's crust

• All ores are minerals
• NOT all minerals are ores

Ore: Minerals from which metals can be extracted profitably and economically

Important Ores:

Gangue (Matrix): Unwanted impurities like sand, clay, and rocky material present in ores

Metallurgy: Extraction of Metals from Ores

Metallurgy is the science of extracting metals from their ores and refining them for use.

Metallurgical Operations

Step 1: Enrichment (Concentration) of Ore

Removal of gangue from ore using various methods:

A. Hydraulic Washing (Gravity Separation/Levigation)

• Based on difference in specific gravity
• Ore mixed with water; lighter gangue washes away
• Heavier ore particles settle
• Used for: Oxide ores of Iron, Tin, Gold, Silver

B. Froth Floatation Process

• Used for sulphide ores (CuFeS₂, PbS, ZnS)
• Finely crushed ore mixed with water and pine oil (~1%)
• Compressed air creates froth
• Ore particles (wetted by oil) rise with froth
• Gangue (wetted by water) settles
• Ore skimmed off from froth

C. Magnetic Separation

• Based on magnetic properties
• Magnetic ore particles attracted by electromagnetic roller
• Non-magnetic gangue falls away
• Used for: Magnetite (iron ore), removing tungsten impurities from tin ore

D. Leaching (Chemical Separation)

• Ore treated with suitable chemical reagents
• Ore dissolves, impurities remain undissolved
• Example - Baeyer's Process for Aluminium:

Al₂O₃ + 2NaOH + H₂O → 2NaAlO₂ + H₂O (Bauxite dissolves, Fe₂O₃ and silicates remain) 2NaAlO₂ + 2H₂O → Al(OH)₃↓ + NaOH (On dilution/CO₂ treatment) 2Al(OH)₃ --1473K-> Al₂O₃ + 3H₂O (On heating)

Step 2: Conversion to Metal Oxide

A. Calcination

• Heating ore in limited air supply below melting point
• Used for carbonate and hydrated oxide ores

Effects:

• Moisture expelled
• Volatile impurities removed
• Carbonate ores → Oxide ores

Examples:

ZnCO₃ → ZnO + CO₂(Calamine to Zinc oxide)Al₂O₃·2H₂O → Al₂O₃ + 2H₂O(Bauxite to Alumina)

B. Roasting

• Heating ore in excess air supply below melting point
• Used for sulphide ores

Effects:

• Moisture expelled
• Volatile impurities removed
• Sulphide ores → Oxide ores

Example:

2ZnS + 3O₂ → 2ZnO + 2SO₂

Both processes carried out in reverberatory furnace

Step 3: Reduction of Metal Oxide

A. For Metals Low in Reactivity Series (Cu, Hg, Ag)

• Heating alone reduces oxides

Examples:

2HgO --Heat--> 2Hg + O₂2Cu₂S + 3O₂ --Heat--> 2Cu₂O + 2SO₂2Cu₂O + Cu₂S --Heat--> 6Cu + SO₂

B. For Metals in Middle of Reactivity Series (Zn, Fe, Pb, Ni, Sn)

(i) Carbon Reduction (Smelting):

ZnO + C → Zn + COSnO₂ + 2C → Sn + 2CO

Carbon monoxide also acts as reducing agent:

ZnO + CO → Zn + CO₂Fe₂O₃ + 3CO → 2Fe + 3CO₂

(ii) Hydrogen Reduction:

• For metals like W, Mo (not reduced by carbon)

WO₃ + 3H₂ → W + 3H₂O

(iii) Aluminothermy:

• For oxides not easily reduced by carbon (Cr₂O₃, MnO₂)
• Aluminium used as reducing agent

Cr₂O₃ + 2Al → Al₂O₃ + 2Cr3MnO₂ + 4Al → 2Al₂O₃ + 3Mn

Thermite Reaction: Used for welding railway tracks

Fe₂O₃ + 2Al → 2Fe (molten) + Al₂O₃

C. For Highly Reactive Metals (Na, Ca, Mg, K, Al)

Electrolytic Reduction:

• Salts electrolysed in molten (fused) state
• Metal deposited at cathode
• Non-metal at anode

Sodium Extraction:

2NaCl --Electrolysis--> 2Na + Cl₂At Cathode: Na⁺ + e⁻ → NaAt Anode: 2Cl⁻ → Cl₂ + 2e⁻

Aluminium Extraction (Hall-Héroult Process):

• Al₂O₃ dissolved in molten cryolite (Na₃AlF₆)
• Graphite electrodes used

At Cathode: Al³⁺ + 3e⁻ → Al(l)At Anode: 2O²⁻ → O₂ + 4e⁻ 2C + O₂ → 2CO 2CO + O₂ → 2CO₂

Step 4: Refining (Purification)

A. Distillation

• For low boiling metals (Zn, Hg, Cd)
• Impure metal heated; pure metal vaporizes
• Vapors condensed to get pure metal

B. Liquation

• For metals with lower M.P. than impurities (Sn, Pb, Bi, Hg)
• Impure metal heated on sloping hearth
• Metal melts and flows down; impurities left behind

C. Oxidation Methods

(i) Bessemerisation: Using Bessemer converter 

(ii) Cupellation: For metals containing impurities that form volatile oxides 

(iii) Poling: For removing oxide impurities from crude metal

Cu₂O + CH₄ → 6Cu + 2H₂O + CO(Green wood poles release hydrocarbons)

D. Electrolytic Refining

• Used for: Cu, Au, Ag, Pb, Zn, Al
• Impure metal = Anode
• Pure metal strip = Cathode
• Electrolyte = Soluble salt of metal

At Anode: M → M^n+ + ne⁻At Cathode: M^n+ + ne⁻ → M (pure)

Anode mud: Less reactive impurities (Au, Ag in copper refining) Solution: More reactive impurities remain dissolved

E. Zone Refining

• For highest purity (Ga, In, Si for semiconductors)
• Based on: Impurities more soluble in molten state
• Moving circular heater melts sections progressively
• Pure metal crystallizes; impurities move to molten zone
• Process repeated until impurities driven to one end

F. Vapor Phase Refining

(i) Mond Process (for Nickel):

Ni (impure) + 4CO -330-350K-> Ni(CO)₄ (vapor)Ni(CO)₄ --Heat~1700K--> Ni (pure) + 4CO

(ii) Van Arkel Method (for Ti, Zr):

Ti (impure) + 2I₂ --500K--> TiI₄ (vapor)TiI₄ --1700K--> Ti (pure) + 2I₂

Flux and Slag

Flux: Substance added during smelting to remove non-fusible impurities Slag: Fusible substance formed when flux combines with impurities

Examples:

SiO₂ + CaO → CaSiO₃ (slag)P₂O₅ + 3CaO → Ca₃(PO₄)₂ (slag)MnO + SiO₂ → MnSiO₃ (slag)

Extraction of Iron and Aluminium

Iron Extraction (Blast Furnace Process)

Raw Materials:

1. Iron ore (Haematite - Fe₂O₃)
2. Coke (carbon)
3. Limestone (CaCO₃) - as flux
4. Hot air

Reactions in Blast Furnace:

Zone 1 (Lower region - 2200K):

C + O₂ → CO₂ + Heat

Zone 2 (Middle region - 1500K):

CO₂ + C → 2CO(Carbon monoxide - reducing agent)

Zone 3 (Upper region - 900K):

Fe₂O₃ + 3CO → 2Fe + 3CO₂(Iron oxide reduced to iron)

Slag Formation:

CaCO₃ → CaO + CO₂CaO + SiO₂ → CaSiO₃ (slag)(Removes silica impurities)

Product: Molten iron collected at bottom; slag floats on top

Types of Iron:

Aluminium Extraction

Ore: Bauxite (Al₂O₃·2H₂O)

Step 1: Baeyer's Process (Purification)

Al₂O₃·2H₂O + 2NaOH → 2NaAlO₂ + 3H₂O2NaAlO₂ + CO₂ + 3H₂O → 2Al(OH)₃ + Na₂CO₃2Al(OH)₃ --Heat 1473K--> Al₂O₃ + 3H₂O

Step 2: Hall-Héroult Process (Electrolysis)

• Pure Al₂O₃ dissolved in molten cryolite (Na₃AlF₆)
• Temperature: ~1000°C
• Graphite anode (consumable)
• Graphite-lined iron cathode

At Cathode: Al³⁺ + 3e⁻ → AlAt Anode: 2O²⁻ → O₂ + 4e⁻ C + O₂ → CO₂

Why Cryolite:

• Lowers melting point of Al₂O₃ (from 2050°C to ~1000°C)
• Increases electrical conductivity
• Reduces energy consumption

Properties of Aluminium:

• Light metal (density 2.7 g/cm³)
• Good conductor of heat and electricity
• Forms protective Al₂O₃ layer (prevents corrosion)
• Highly reactive but protected by oxide layer

Uses of Aluminium:

1. Aircraft construction (duralumin alloy)
2. Electrical transmission wires
3. Food packaging (foils)
4. Cooking utensils
5. Alloys (duralumin, magnalium)

Corrosion of Metals

Understanding Corrosion

Corrosion is the slow destruction of metals due to chemical reactions with oxygen, moisture, CO₂, SO₂, H₂S, and other atmospheric substances. It weakens metals progressively and can lead to complete destruction.

Mechanism of Corrosion

Electrochemical Process:

M → M⁺ + e⁻ (Metal loses electrons)

• Electrons flow in presence of moisture (electrolyte)
• Metal ions move from one part to another
• Oxidation occurs at one site, reduction at another
• Forms an electrochemical cell

Rusting of Iron

Rust = Hydrated iron(III) oxide [Fe₂O₃·xH₂O + Fe(OH)₃]

Essential Conditions:

1. Presence of oxygen/air
2. Presence of water/moisture

Reaction:

Fe + H₂O + O₂ → Fe₂O₃·xH₂O + Fe(OH)₃(Iron) (Rust - reddish brown, flaky)

Factors Accelerating Corrosion:

1. Contact with Different Metals:
   • Forms galvanic cell
   • More reactive metal corrodes faster
   • Example: Iron corrodes faster when in contact with copper
2. Atmospheric Pollutants:
   • CO₂, SO₂, H₂S accelerate corrosion
   • Industrial areas have higher corrosion rates
3. Moisture and Humidity:
   • Higher humidity increases corrosion
   • Coastal areas experience faster corrosion due to salt
4. Acidic Environment:
   • Acids dissolve protective oxide layers

Examples of Corrosion:

• Silver: Forms black Ag₂S (silver sulphide) with atmospheric sulphur
• Copper: Forms green basic copper carbonate [CuCO₃·Cu(OH)₂]
• Iron: Forms reddish-brown rust

Prevention of Corrosion

1. Barrier Protection

Painting:

• Paint forms protective layer
• Prevents contact with air and moisture
• Paints with red lead, zinc chromate, lead chromate most effective

Greasing/Oiling:

• For machinery parts
• Temporary protection
• Needs regular application

Coating with Enamel:

• Provides glossy protective layer
• Used for kitchen utensils

2. Galvanization

• Coating iron with zinc layer
• Zinc more reactive than iron
• Forms protective basic zinc carbonate [Zn(OH)₂·ZnCO₃]
• Used for: Water pipes, buckets, roofing sheets

Process: Iron article dipped in molten zincResult: Zinc coating prevents rusting

3. Electroplating

• Coating with less reactive, corrosion-resistant metals
• Chromium, Nickel, Tin, Aluminium commonly used
• Provides attractive appearance
• Used for: Automobile parts, taps, handles

Process:

• Metal to be coated = Cathode
• Coating metal = Anode
• Electrolyte = Salt of coating metal

4. Alloying

Stainless Steel:

• Iron + Chromium (18%) + Nickel (8%) + Carbon (0.08%)
• Highly resistant to corrosion
• Used for: Cutlery, surgical instruments, construction

5. Coating with Metal Oxide

• Iron objects coated with Fe₃O₄
• Forms protective black oxide layer

6. Sacrificial Protection

• More reactive metal connected to iron structure
• Reactive metal oxidizes preferentially
• Magnesium or Zinc strips used
• Used for: Underground iron pipes, ship hulls

Example: Mg/Zn connected to Fe pipe Mg/Zn corrodes; Fe protected

7. Anodising (for Aluminium)

• Forms thick Al₂O₃ layer
• Aluminium article made anode
• Electrolysed with dilute H₂SO₄
• Oxygen at anode reacts to form thick oxide
• Oxide layer can be dyed for decorative finish

Applications: Aircraft bodies, window frames, cookware

8. Chemical Coating

• Bisphenol, oxides applied to metal surface
• Concrete coating for large structures
• Phosphate coating (Fe/Mn phosphates) for atmospheric corrosion
• Glycol in coolants for automobiles

9. Preventive Measures

For Copper/Brass:

• p-Chlorobenzohydroxamic acid coating

For Automobiles:

• Cyanovanadate in radiators

For Steel/Nuclear Plants:

• Anodic potential method

Alloys: Homogeneous Metal Mixtures

What are Alloys?

Alloy: Homogeneous mixture of two or more metals, or a metal with a non-metal, where the major component is a metal.

Purpose of Alloying:

• Pure metals often too soft, brittle, or chemically reactive
• Alloys have improved properties:
  • Increased hardness
  • Enhanced corrosion resistance
  • Better color and luster
  • Modified melting points
  • Reduced brittleness

Amalgam: Special alloy containing mercury as one component

Preparation of Alloys

Methods:

1. Melting metals in proper proportions; cooling and solidifying the melt
2. Compressing constituents under high pressure
3. Simultaneous electrodeposition of metals

Properties of Alloys

1. Harder than constituent metals but less ductile/malleable
2. Resistant to corrosion
3. Melting point may be higher or lower than constituents
4. Improved properties compared to pure metals
5. Different color from constituent metals

Example: Silver and zinc (white) form pink alloy

Types of Alloys

1. Ferrous Alloys - Contain iron

• Steel, Stainless steel, Nickel steel

2. Non-Ferrous Alloys - Do not contain iron

• Bronze, Brass, Duralumin

Important Alloys and Their Uses

Alloys of Aluminium

Alloys of Copper

Alloys of Silver

Alloys of Tin and Lead

Alloy Steels (Iron Alloys)

Steel: Alloy of iron and carbon (0.15-1.7% carbon)

Note: Electrical conductivity of alloys is generally less than pure metals.

Special Alloys

Gold Alloys:

• 24 carat gold = Pure gold (100%)
• 22 carat gold = 22 parts gold + 2 parts Cu/Ag (91.67% gold)
• 18 carat gold = 18 parts gold + 6 parts Cu/Ag (75% gold)

Pure gold too soft for jewelry; alloyed with copper/silver for hardness.

Magnetic Metals: Only iron, cobalt, nickel are naturally magnetic. Steel (mostly iron) is also magnetic.

Strategic Metals: Titanium, chromium, manganese, zirconium

• Used in atomic energy, space science, jet engines, defence equipment

Extreme Properties:

• Heaviest metal: Iridium
• Lightest metal: Lithium

Important Metal Compounds

1. Silver Nitrate (AgNO₃) - Lunar Caustic

Properties:

• Colorless, transparent crystals
• Soluble in water

Uses:

• Laboratory reagent
• Marking inks
• Photography
• Silvering of mirrors
• Manufacturing other silver salts

2. Silver Bromide (AgBr)

Properties:

• Pale yellow crystals
• Insoluble in water
• Light sensitive

Uses: Photography (light-sensitive films)

3. Potash Alum [K₂SO₄·Al₂(SO₄)₃·24H₂O]

Properties:

• Double salt of aluminium and potassium sulphate
• Soluble in water
• Contains large amount of crystalline water
• Expands on heating (loses water)

Uses: Water softening

4. Copper Sulphate (CuSO₄·5H₂O) - Blue Vitriol

Properties:

• Blue crystalline substance
• Releases water on heating

CuSO₄·5H₂O --373K--> CuSO₄·H₂O --413K--> CuSO₄ --443K--> CuO + SO₂

Uses:

• Electroplating
• Electric batteries
• Bordeaux mixture (CuSO₄ + lime) - fungicide
• Dyeing clothes
• Wood preservation

Important Non-Metal Elements

Sulphur (S)

Atomic Number: 16
Atomic Mass: 32
Electronic Configuration: 1s² 2s² 2p⁶ 3s² 3p⁴
Valencies: +2, +4, +6, -2

Occurrence:

• Free state: Volcanic regions
• Combined state:
  • Sulphates: Gypsum (CaSO₄·2H₂O), Epsom salt (MgSO₄·7H₂O), Glauber's salt (Na₂SO₄·10H₂O)
  • Sulphides: Cinnabar (HgS), Galena (PbS), Iron pyrites (FeS₂), Zinc blende (ZnS)
  • Organic matter: Insulin, natural gas, petroleum, coal

Allotropy:

A. Crystalline Allotropes:

Transition: Rhombic ⇌ Monoclinic at 368.6K (transition temperature)

B. Non-Crystalline Allotropes:

1. Plastic Sulphur (γ-sulphur):
   • Obtained by pouring boiling sulphur in cold water
   • Rubber-like, soft substance
   • Chain structure (S∞)
   • Unstable; converts to rhombic sulphur
2. Milk of Sulphur (δ-sulphur):
   • White colored
   • Used in medicines
3. Colloidal Sulphur:
   • Prepared by: H₂S + 2HNO₃ → 2NO₂ + 2H₂O + S
   • Soluble in CS₂
   • Used in medicines

Uses:

• Manufacturing sulphuric acid
• Gunpowder and match sticks
• Insecticide
• Medicines (skin diseases, blood purification)
• Vulcanization of rubber

Phosphorus (P)

Atomic Number: 15
Atomic Mass: 31
Electronic Configuration: 1s² 2s² 2p⁶ 3s² 3p³
Valencies: +3, +5

Name Origin: Greek - phos (light) + phero (carry) = "light carrier"

Occurrence:

• Combined state only (highly reactive)
• Minerals:
  • Phosphorite [Ca₃(PO₄)₂]
  • Fluorapatite [3Ca₃(PO₄)₂·CaF₂]
  • Chlorapatite [3Ca₃(PO₄)₂·CaCl₂]

Allotropes:

Uses:

• White phosphorus: Smoke clouds, fireballs, colored match sticks
• Red phosphorus: Safety match sticks, phosphor bronze alloy
• Rat poisons: Zinc phosphide, calcium phosphide

Important Non-Metal Compounds

1. Ozone (O₃)

Structure: Bent molecule (allotrope of oxygen)

Properties:

• Strong oxidizing agent
• Provides nascent oxygen

Uses:

• Artificial silk production
• Bleaching agent
• Disinfectant
• Purifying drinking water and air
• Manufacturing potassium permanganate

Environmental Importance: Ozone layer in stratosphere blocks harmful UV radiation from sun

2. Hydrogen Peroxide (H₂O₂)

Structure: H-O-O-H (non-linear)

Properties:

• Both oxidizing and reducing agent
• Oxidizing properties predominant

Uses:

• Rejuvenating old paintings (converts PbS to PbSO₄)

PbS + 4H₂O₂ → PbSO₄ + 4H₂O (Yellow) → (White)

• Disinfectant for wounds (dilute solution)
• Rocket fuel oxidant (concentrated)
• Bleaching: Silk, hair, ivory, wool, wood
• Medicines and cosmetics

3. Ammonia (NH₃)

Structure: Pyramidal (N with 3 H atoms)

Uses:

• Manufacturing nitrogen fertilizers (ammonium nitrate, urea, ammonium phosphate, ammonium sulphate)
• Manufacturing nitric acid
• Refrigerant/freezing agent
• Synthetic silk production
• Explosive manufacturing
• Removing oil/grease stains
• Medicines and cosmetics (dilute)

4. Nitric Acid (HNO₃)

Structure: O=N-OH with double-bonded O

Properties: Strong oxidizing agent

Uses:

• Manufacturing ammonium nitrate (fertilizer)
• Manufacturing explosives (potassium nitrate, nitroglycerin, TNT)
• Rocket fuel oxidation
• Dyes, medicines, artificial silk
• Purifying gold and silver
• Manufacturing sulphuric acid

5. Sulphuric Acid (H₂SO₄)

Structure: O=S(=O)(OH)₂

Properties: Strong acid, called "King of Acids"

Uses:

• Manufacturing HCl, HNO₃, fertilizers
• Petroleum refining
• Dye industry
• Metallurgy (metal extraction)
• Batteries
• Electroplating
• Detergent industry
• Hundreds of industrial applications

6. Hydrochloric Acid (HCl)

Common Name: Acid of common salt

Uses:

• Laboratory reagent
• Aqua Regia (3 HCl : 1 HNO₃) - dissolves noble metals like gold, platinum
• Steel and iron industry
• Textile industry
• Manufacturing gum and dyes
• Producing chlorine gas (for bleaching powder, water purification)
• Synthesis of metal and non-metal chlorides

7. Ammonium Chloride (NH₄Cl)

Common Name: Nausadar

Uses:

• Soldering material
• Polishing utensils
• Electric batteries
• Dyeing clothes
• Medicine
• Laboratory reagent

8. Silica (SiO₂)

Occurrence: Sand (most abundant form)

Uses:

• Glass manufacturing
• Jasper (semi-precious stone)
• Mother silicates (electrical appliances, laboratory equipment, optical lenses, prisms)
• Silicon production (for semiconductors, transistors, alloys)
• Kieselguhr (allotrope) - dynamite manufacturing

Chemical Formulas

Conclusion

Learn metals and non-metals is fundamental to chemistry and has immense practical significance. From the extraction of iron and aluminium that forms the backbone of modern infrastructure to preventing corrosion that saves billions in maintenance costs, these concepts directly impact our daily lives.

Chapter Important Notes:

1. Physical properties clearly distinguish metals from non-metals
2. Chemical reactivity varies systematically as shown in the reactivity series
3. Extraction methods depend on the metal's position in the reactivity series
4. Corrosion prevention requires understanding of electrochemical processes
5. Alloys combine the best properties of constituent metals for practical applications

Learn these concepts provides a strong foundation for advanced chemistry and practical applications in metallurgy, materials science, and engineering.