Metals and Non-Metals – Notes

1. Introduction

There are 115 known chemical elements.
Elements show similarities and differences in their properties.
On the basis of properties, elements are classified into:
(a) Metals
(b) Non-Metals

2. Important Terms

Malleable → can be hammered into thin sheets.
Ductile → can be drawn into thin wires.
Brittle → breaks easily when hammered or stretched.

3. Metals

Definition:

Metals are elements that:
- Conduct heat and electricity
- Are malleable and ductile
- Form positive ions by losing electrons (electropositive).

General Properties of Metals:

Good conductors of heat and electricity
Malleable and ductile
Lustrous (shiny)
Hard, strong, heavy
Sonorous (produce ringing sound)
Mostly solids (except mercury which is liquid)

Examples: Iron, Aluminium, Copper, Silver, Gold, Platinum, Zinc, Sodium, Potassium, Calcium, Magnesium.

Abundance in Earth’s Crust:

Aluminium (7%) → most abundant metal
Iron (4%)
Others: Calcium, Sodium, Potassium, Magnesium

Uses of Metals:

Utensils, vehicles, machinery, wires, ornaments, construction, alloys.
Important for national economy.

4. Non-Metals

Definition:

Non-metals are elements that:
- Do not conduct heat and electricity
- Are not malleable or ductile (brittle)
- Form negative ions by gaining electrons (electronegative)
- Exception: Hydrogen forms H⁺ ions by losing electrons.

General Properties of Non-Metals:

Poor conductors of heat and electricity
Brittle, non-malleable, non-ductile
Dull (not shiny)
Generally soft and light
Non-sonorous (do not produce ringing sound)
Exist as solids, liquids, or gases
- Bromine → only liquid non-metal

Examples: Carbon, Sulphur, Phosphorus, Silicon, Hydrogen, Oxygen, Nitrogen, Chlorine, Bromine, Iodine, Helium, Neon, Argon.

Abundance in Earth’s Crust:

Oxygen (50%) → most abundant non-metal
Silicon (26%)
Others: Phosphorus, Sulphur

Importance of Non-Metals:

Carbon → basis of life (proteins, fats, carbohydrates, vitamins, enzymes)
Oxygen → essential for breathing & combustion
Nitrogen → controls combustion rate, used in fertilizers
Sulphur → present in hair, onion, wool; used in fungicides, explosives
Hydrogen + Oxygen → form water (oceans)
Chlorine → found in oceans as salts

5. Key Difference Between Metals & Non-Metals

Property	Metals	Non-Metals
Conductivity	Good conductor	Poor conductor
Malleability & Ductility	Malleable & ductile	Brittle
Appearance	Lustrous (shiny)	Dull
Strength	Hard & strong	Soft & weak
Sound	Sonorous (ringing sound)	Non-sonorous
Ions in Reactions	Form positive ions (lose electrons)	Form negative ions (gain electrons)
State at Room Temperature	Mostly solids (except mercury – liquid)	Solids, gases, and one liquid (bromine)

✅ Quick Remember Tip:

Metals → Electropositive (lose e⁻, form cations)
Non-Metals → Electronegative (gain e⁻, form anions)

Physical Properties of Metals

Metals show certain characteristic physical properties that make them different from non-metals.

1. Malleability

Metals can be hammered into thin sheets without breaking.
This property is called malleability.
Examples:
- Gold & Silver → best malleable (used for foils, jewellery).
- Aluminium foils → used for food packing, bottle caps.
- Copper & Iron sheets → used for utensils, buckets, tanks.

2. Ductility

Metals can be drawn into thin wires.
This property is called ductility.
Examples:
- Gold → most ductile (1 g gold → 2 km long wire).
- Silver, Copper, Aluminium → very ductile (used in electric wires).
- Tungsten → wires used in bulb filaments.

3. Good Conductors of Heat

Metals allow heat to pass through them easily → thermal conductivity.
Best conductor: Silver
Copper & Aluminium → very good conductors (used in cooking utensils, boilers).
Poor conductors: Lead, Mercury
Reason: Free-moving electrons carry heat from hot end to cold end.

4. Good Conductors of Electricity

Metals allow electricity (electric current) to pass through them.
Best conductor: Silver → followed by Copper, Gold, Aluminium.
Copper & Aluminium → used in electric wires.
Poor conductors: Iron, Mercury
Reason: Free electrons move easily in metals and carry current.
Note: Electric wires are covered with insulating plastic (like PVC) to prevent shocks.

5. Lustre (Shiny Appearance)

Freshly cut metals have a shiny surface → metallic lustre.
Gold, Silver, Copper → highly lustrous (used in jewellery, mirrors).
Silver → excellent reflector of light (used in mirrors).
Metals lose shine over time due to corrosion (layer of oxide/carbonate/sulphide forms). Shine can be restored by polishing.

6. Hardness

Most metals are hard and cannot be cut with a knife.
Exceptions: Sodium & Potassium → soft metals, can be cut with a knife.
Iron, Copper, Aluminium → very hard.

7. Strength

Metals are generally strong and can bear heavy loads.
Example: Iron (in the form of steel) → used in bridges, buildings, machines, vehicles.
Exceptions: Sodium & Potassium → weak metals.

8. State at Room Temperature

Metals are generally solids at room temperature.
Exception: Mercury → only liquid metal.

9. Melting and Boiling Points

Metals generally have high melting & boiling points.
Examples:
- Iron → MP = 1535°C
- Copper → MP = 1083°C
Exceptions:
- Sodium (98°C), Potassium (64°C), Gallium (30°C), Cesium (28°C) → low melting points.
- Gallium & Cesium → melt in hand!

10. Density

Metals are generally heavy substances (high density).
Example: Iron → 7.8 g/cm³
Exceptions: Sodium (0.97 g/cm³) & Potassium (0.86 g/cm³) → very light metals.

11. Sonority

Metals produce a ringing sound when struck → called sonorousness (sonority).
Used in making bells, gongs, musical instruments (sitar, violin strings).

12. Colour

Most metals are silver-grey in colour.
Exceptions:
- Copper → reddish-brown
- Gold → yellow

✅ Summary – Key Properties of Metals

Malleable (sheets), Ductile (wires), Lustrous (shiny), Sonorous (sound).
Good conductors of heat & electricity.
Hard, strong, high density, high MP/BP (with exceptions).
Usually silver-grey, except copper (red) & gold (yellow).
All solids (except mercury – liquid).

Physical Properties of Non-Metals

Non-metals generally show properties opposite to metals.

1. Non-metals are not malleable or ductile (they are brittle)

Non-metals break into pieces when hammered or stretched.
Cannot be made into thin sheets or wires.
Property is called brittleness.
Examples: Sulphur, Phosphorus, Carbon (all brittle).
⚠️ Note: Brittleness is considered only for solid non-metals (not liquids or gases).

2. Poor Conductors of Heat and Electricity

Non-metals lack free electrons → cannot conduct heat/electricity.
Examples: Sulphur, Phosphorus (do not conduct).
Exception: Graphite (allotrope of carbon) → good conductor of electricity, used in electrodes.

3. Not Lustrous (Dull Appearance)

Most non-metals are dull (no metallic shine).
Examples: Sulphur, Phosphorus.
Exception: Iodine → shiny (lustrous).

4. Generally Soft

Solid non-metals are usually soft.
Examples: Sulphur, Phosphorus.
Exception: Diamond (allotrope of carbon) → hardest natural substance known.

5. Not Strong

Non-metals have low strength, break easily under pressure.
Example: Graphite sheets break easily.

6. Different Physical States at Room Temperature

Non-metals can exist as solids, liquids, or gases.
Solids: Carbon, Sulphur, Phosphorus
Liquid: Bromine (only liquid non-metal)
Gases: Hydrogen, Oxygen, Nitrogen, Chlorine

7. Low Melting and Boiling Points

Most non-metals melt/boil at low temperatures.
Example: Sulphur → MP = 115°C.
Exception: Diamond → very high MP (> 3500°C).

8. Low Densities

Non-metals are usually light substances.
Example: Sulphur → density = 2 g/cm³.

9. Non-Sonorous

Do not produce a ringing sound when struck.

10. Different Colours

Non-metals exist in variety of colours:
- Sulphur → yellow
- Phosphorus → white/red
- Graphite → black
- Chlorine → yellowish-green
- Oxygen & Hydrogen → colourless

Important Exceptions

Some non-metals show properties like metals, and some metals show properties like non-metals:

Conductivity → Graphite (non-metal) conducts electricity.
Lustre → Iodine (non-metal) is lustrous.
Hardness → Diamond (non-metal) is extremely hard.
Melting/Boiling Point → Diamond has very high MP/BP like metals.
Soft Metals → Sodium, Potassium (metals) are soft like non-metals.
Low Density Metals → Lithium, Sodium, Potassium (metals) are light like non-metals.
Liquid Metal → Mercury (metal) is liquid at room temperature.
Low MP Metals → Sodium, Potassium, Gallium, Cesium (metals) have low melting points like non-metals.

✅ Summary – Physical Properties of Non-Metals

Brittle, not malleable or ductile.
Poor conductors (except graphite).
Dull (except iodine).
Soft (except diamond).
Low strength, low density, low MP/BP (except diamond).
Exist in solid, liquid, and gas states.
Non-sonorous, many colours.

👉

CHEMICAL PROPERTIES OF METALS

Metals show different chemical behaviours. Their main chemical properties are explained below:

1. Reaction of Metals with Oxygen

Metals react with oxygen (O₂) of air to form metal oxides.


\text{Metal + Oxygen → Metal Oxide}

Most metal oxides are basic in nature. They turn red litmus → blue.
Some oxides dissolve in water to form alkalis.
A few oxides (like Al₂O₃ and ZnO) are amphoteric (react with both acids and bases).

Examples:

(i) Highly Reactive Metals (Na, K, Li)

React with oxygen at room temperature.
Catch fire if exposed to air → hence kept in kerosene oil.


4Na + O₂ \; \to \; 2Na₂O

2K + O₂ ; \to ; K₂O 

These oxides dissolve in water to form alkalis:


Na₂O + H₂O \to 2NaOH

K₂O + H₂O \to 2KOH 

(ii) Magnesium (Mg)

Does not react at room temperature.
On heating, burns with bright white light to form MgO.


2Mg + O₂ \; \to \; 2MgO

MgO is basic, partly soluble in water:


MgO + H₂O \to Mg(OH)₂

(iii) Aluminium (Al)

Burns on heating to form aluminium oxide (Al₂O₃).


4Al + 3O₂ \to 2Al₂O₃

Al₂O₃ is amphoteric:
(a) With acids → behaves like a base:


Al₂O₃ + 6HCl \to 2AlCl₃ + 3H₂O


Al₂O₃ + 2NaOH \to 2NaAlO₂ + H₂O

(iv) Zinc (Zn)

Burns only on strong heating → forms ZnO (amphoteric).


2Zn + O₂ \to 2ZnO

Reactions of ZnO:
(a) With acids:


ZnO + 2HCl \to ZnCl₂ + H₂O


ZnO + 2NaOH \to Na₂ZnO₂ + H₂O

(v) Iron (Fe)

Does not burn, but on heating:


3Fe + 2O₂ \to Fe₃O₄

(Iron filings burn, but a big piece does not.)

(vi) Copper (Cu)

Does not burn in air.
On prolonged heating:


2Cu + O₂ \to 2CuO

(vii) Silver (Ag) & Gold (Au)

Do not react with oxygen even at high temperature.
They are the least reactive metals.

Key Points to Remember

Most metal oxides are basic, but Al₂O₃ and ZnO are amphoteric.
Na, K, Li react vigorously with O₂ (very reactive).
Mg burns with a white flame, forming MgO (basic).
Al & Zn oxides react with both acids and bases.
Fe, Cu react slowly with oxygen.
Ag & Au do not react → noble metals.

✨ Exam Tip: Always write balanced chemical equations with states (s, l, g, aq) and note whether the oxide is basic, amphoteric, or acidic.

Reaction of Metals with Water

Metals react with water to form metal hydroxides (or oxides) and hydrogen gas.
The nature and intensity of reaction depend on the reactivity of the metal.

General Reactions

With cold/hot water:


\text{Metal + Water → Metal Hydroxide + Hydrogen}

With steam (very hot water vapour):


\text{Metal + Steam → Metal Oxide + Hydrogen}

Examples of Reactions

(i) Potassium (K)

Reacts violently with cold water.
Reaction is highly exothermic → hydrogen catches fire immediately.


2K + 2H₂O \;\to\; 2KOH + H₂ ↑ + \text{Heat}

(ii) Sodium (Na)

Reacts vigorously with cold water.
Hydrogen burns with little explosions.


2Na + 2H₂O \;\to\; 2NaOH + H₂ ↑ + \text{Heat}

(iii) Calcium (Ca)

Reacts with cold water but less violently.
Hydrogen formed does not catch fire (reaction less exothermic).
Calcium floats because bubbles of H₂ stick to its surface.


Ca + 2H₂O \;\to\; Ca(OH)₂ + H₂ ↑

(iv) Magnesium (Mg)

No reaction with cold water.
With hot water:


Mg + 2H₂O \;\to\; Mg(OH)₂ + H₂ ↑


Mg + H₂O (g) \;\to\; MgO + H₂ ↑

(v) Aluminium (Al)

No reaction with cold/hot water (protected by thin Al₂O₃ layer).
With steam:


2Al + 3H₂O (g) \;\to\; Al₂O₃ + 3H₂ ↑

(vi) Zinc (Zn)

No reaction with cold/hot water.
With steam:


Zn + H₂O (g) \;\to\; ZnO + H₂ ↑

(vii) Iron (Fe)

No reaction with cold/hot water.
With steam (red hot iron):


3Fe + 4H₂O (g) \;\to\; Fe₃O₄ + 4H₂ ↑

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

Do not react with water or steam (very unreactive).

Order of Reactivity with Steam


\text{Magnesium > Aluminium > Zinc > Iron}

Why Metals Displace Hydrogen from Water

Water is slightly ionised:


H₂O \;\rightleftharpoons\; H⁺ + OH⁻

Less reactive metals (like Cu, Ag, Au) cannot give electrons easily → no reaction.
Rule: Only metals above hydrogen in the reactivity series displace H₂ from water.

Summary Table: Reaction of Metals with Water

Metal	Cold Water	Hot Water	Steam	Nature of Reaction
Potassium	Violent, H₂ catches fire	—	—	Very reactive
Sodium	Vigorous, H₂ catches fire	—	—	Very reactive
Calcium	Slow, less violent, H₂ not burnt	—	—	Reactive
Magnesium	No reaction	Slow	Vigorous	Moderately reactive
Aluminium	No reaction	No reaction	Forms Al₂O₃ + H₂	Reactive (protected by oxide layer)
Zinc	No reaction	No reaction	Forms ZnO + H₂	Reactive
Iron	No reaction	No reaction	Forms Fe₃O₄ + H₂	Less reactive
Pb, Cu, Ag, Au	No reaction	No reaction	No reaction	Unreactive

✅ Quick Recap for Exams:

K & Na: violent with cold water.
Ca: less violent, floats on water.
Mg: reacts with hot water, faster with steam.
Al, Zn, Fe: react only with steam.
Cu, Ag, Au: no reaction.

🌟 Reaction of Metals with Dilute Acids

🔹 General Rule

Most metals react with dilute acids (like HCl or H₂SO₄) to form:


  \text{Metal + Dilute acid → Metal salt + Hydrogen gas}

Reaction depends on metal reactivity.

1️⃣ Reaction with Dilute Hydrochloric Acid (HCl)

Metals + HCl → Metal chloride + H₂ gas

Examples:

Sodium (very reactive, violent reaction):


   2Na + 2HCl → 2NaCl + H₂ ↑

Magnesium (rapid reaction):


   Mg + 2HCl → MgCl₂ + H₂ ↑

Aluminium (slow at first due to oxide layer, then rapid):


   2Al + 6HCl → 2AlCl₃ + 3H₂ ↑

Zinc (moderate reaction):


   Zn + 2HCl → ZnCl₂ + H₂ ↑

Iron (slow reaction):


   Fe + 2HCl → FeCl₂ + H₂ ↑

Copper (no reaction):


   Cu + HCl → \text{No reaction}

2️⃣ Reactivity Order with Dilute HCl

Experiment shows bubble formation rate:


  \text{Mg > Al > Zn > Fe > Cu (no reaction)}

Key Point: Metals above hydrogen in reactivity series displace hydrogen from acids.
Metals below hydrogen (like Cu, Ag, Au) do not.

3️⃣ Reaction with Dilute Sulphuric Acid (H₂SO₄)

Metals + H₂SO₄ → Metal sulphate + H₂ gas

Examples:


2Na + H₂SO₄ → Na₂SO₄ + H₂ ↑

Mg + H₂SO₄ → MgSO₄ + H₂ ↑ 


2Al + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂ ↑

Zn + H₂SO₄ → ZnSO₄ + H₂ ↑ 


Fe + H₂SO₄ → FeSO₄ + H₂ ↑

Cu + H₂SO₄ → \text{No reaction} 

4️⃣ Reaction with Dilute Nitric Acid (HNO₃)

Normally: No H₂ gas is evolved ❌
- Because HNO₃ is a strong oxidising agent.
- It oxidises H₂ → H₂O.
- Nitrates + nitrogen oxides (N₂O, NO, NO₂) are formed.
Exception (with very dilute HNO₃):
Hydrogen gas evolves with Mg and Mn.

Examples:


Mg + 2HNO₃ (very dilute) → Mg(NO₃)₂ + H₂ ↑

Mn + 2HNO₃ (very dilute) → Mn(NO₃)₂ + H₂ ↑ 

5️⃣ Key Concept – Why Some Metals Displace Hydrogen?

Acids provide H⁺ ions.
Reactive metals lose electrons easily → reduce H⁺ → H₂ gas.
Less reactive metals (Cu, Ag, Au) cannot lose electrons easily → no reaction.

6️⃣ Summary Table: Reactivity with Dilute Acids

Metal	Reaction with Dilute Acids	Observation
Na, K	Very violent, explosive	Salt + H₂ ↑
Mg	Rapid	Salt + H₂ ↑
Al	Slow (oxide layer) → Rapid	Salt + H₂ ↑
Zn	Moderate	Salt + H₂ ↑
Fe	Slow	Salt + H₂ ↑
Cu, Ag, Au	No reaction	–

✅ In short:

Metals above hydrogen in activity series → displace H from acids.
Reactivity order with acids: Na > Mg > Al > Zn > Fe > Cu (no reaction).
With HNO₃: Usually no H₂, except very dilute acid with Mg/Mn.

🌟 Aqua Regia

🔹 Definition

Aqua regia is a freshly prepared mixture of:
- 1 part concentrated nitric acid (HNO₃)
- 3 parts concentrated hydrochloric acid (HCl)
Ratio = 1 : 3 (HNO₃ : HCl)

🔹 Properties

Highly corrosive and fuming liquid.
Known as "Royal Water" because it can dissolve noble metals (gold & platinum) which neither HCl nor HNO₃ alone can dissolve.

🔹 Why Can Aqua Regia Dissolve Gold & Platinum?

Conc. HNO₃ (oxidising agent) produces nascent chlorine from HCl.
Nascent chlorine ([Cl]) + nitrosyl chloride (NOCl) help in dissolving noble metals.

Reactions Involved:


HNO₃ + 3HCl → NOCl + 2H₂O + Cl₂

The released Cl₂ and nascent chlorine attack Au and Pt to form soluble chloro-complexes.


Au + 3Cl₂ → AuCl₃

Pt + 4Cl₂ → PtCl₄ 

🔹 Uses of Aqua Regia

To dissolve gold and platinum.
Used in purification and refining of precious metals.
Used in laboratories for cleaning glassware (removes stubborn metal stains).

🔹 Key Points to Remember

Freshly prepared (unstable if stored).
Ratio = 1 part conc. HNO₃ : 3 parts conc. HCl.
Strong oxidising + chlorinating mixture.
Can dissolve noble metals which resist other acids.

✅ In short:
Aqua regia is a powerful mixture of nitric acid & hydrochloric acid (1:3), capable of dissolving even noble metals like gold and platinum due to the combined action of oxidising nitric acid and nascent chlorine from hydrochloric acid.

⚡ The Reactivity Series of Metals (Activity Series)

🔹 Definition

The reactivity series (or activity series) is the arrangement of metals in decreasing order of their reactivity.
Most reactive metals are placed at the top, while the least reactive metals are placed at the bottom.
Reactivity is decided based on reactions with oxygen, water, acids, and displacement reactions.

🔹 Reactivity Series (Important Metals)

Most reactive → Least reactive

Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Tin (Sn)
Lead (Pb)
[Hydrogen (H)]
Copper (Cu)
Mercury (Hg)
Silver (Ag)
Gold (Au)

🔹 Key Facts

Potassium (K) = Most reactive (top of series).
Gold (Au) = Least reactive (bottom of series).
Hydrogen (H) = Not a metal, but placed in the series since it behaves like a metal (forms H⁺ ions).
Metals at the top are found in combined state (ores).
Metals at the bottom (like Ag, Au) are found in native (free) state due to low reactivity.

🔹 Why Some Metals are More Reactive?

Metals react by losing electrons to form positive ions (M → M⁺ + e⁻).
If a metal loses electrons easily, it is highly reactive (e.g., Na, K).
If a metal loses electrons slowly, it is less reactive (e.g., Fe, Cu).

🔹 Metals More Reactive than Hydrogen

K, Na, Ca, Mg, Al, Zn, Fe, Sn, Pb
These metals can displace hydrogen from water and acids, producing H₂ gas.

🔹 Metals Less Reactive than Hydrogen

Cu, Hg, Ag, Au
These metals cannot displace hydrogen from water or acids.

🔹 Importance of Reactivity Series

✅ Helps predict whether a displacement reaction will occur.
✅ Explains reactions of metals with oxygen, water, and acids.
✅ Helps understand extraction of metals from ores (more reactive metals require stronger reduction methods).

🔹 In Short:

Above H → Metal reacts with water/acids → H₂ gas released.
Below H → No reaction with water/acids → No H₂ gas.

👉

⚡ Reaction of Metals with Salt Solutions

🔹 General Rule

When a more reactive metal is placed in the salt solution of a less reactive metal,
→ the more reactive metal displaces the less reactive one from its salt.
Reaction (general form):
Salt solution of Metal B + Metal A → Salt solution of Metal A + Metal B
(Only if A is more reactive than B).

This is called a displacement reaction.

🔹 Examples of Displacement Reactions

1. Zinc + Copper Sulphate

When zinc strip is dipped in blue CuSO₄ solution:
CuSO₄ (aq) + Zn (s) → ZnSO₄ (aq) + Cu (s)
Observations:
- Blue colour fades → colourless solution (ZnSO₄).
- Red-brown copper is deposited on zinc.
Reason: Zinc is more reactive than copper.

2. Iron + Copper Sulphate

When an iron nail is dipped in CuSO₄ solution:
CuSO₄ (aq) + Fe (s) → FeSO₄ (aq) + Cu (s)
Observations:
- Blue solution turns green (FeSO₄).
- Red-brown copper deposits on iron nail.
Reason: Iron is more reactive than copper.

3. Copper + Silver Nitrate

When copper strip is placed in AgNO₃ solution:
2AgNO₃ (aq) + Cu (s) → Cu(NO₃)₂ (aq) + 2Ag (s)
Observations:
- Colourless solution turns blue (Cu(NO₃)₂).
- Grey-white silver is deposited on copper.
Reason: Copper is more reactive than silver.

🔹 Important Points

If a less reactive metal is placed in the salt solution of a more reactive metal → No reaction occurs.
Example: Cu in ZnSO₄ solution → ❌ No reaction.

🔹 Applications

Reactivity Series helps predict whether displacement will occur.
Extraction of metals – highly reactive metals displace less reactive metals from their compounds.
Explains corrosion of metals (iron vessels corrode in CuSO₄ solution).

🔹 NCERT Sample Problems (Key Takeaways)

Q1. Cu in AgNO₃ → Cu displaces Ag → Copper is more reactive than silver.
Q2. CuSO₄ in iron pot → Fe displaces Cu → pot develops holes.
Q3. Zn + FeSO₄ → Zn displaces Fe → Colour changes from green to colourless.
Q4. Displacement occurs only if the solid metal is above the other metal in the reactivity series.

🔹 Golden Rule (to remember):

👉 A metal can displace another metal only if it is higher in the reactivity series.

✅

Here are clear, exam-ready notes on Reaction of Metals with Chlorine for students 👇

⚡ Reaction of Metals with Chlorine

🔹 General Rule

Metals react with chlorine (Cl₂) to form ionic chlorides.
In this reaction:
- Metal atoms lose electrons → form positively charged metal ions (M⁺ / M²⁺ / M³⁺, etc.).
- Chlorine atoms gain electrons → form negatively charged chloride ions (Cl⁻).
These oppositely charged ions are held together by electrostatic forces → forming ionic compounds.

👉 Equation (general):
Metal (M) + Chlorine (Cl₂) → Metal chloride (MClₓ)

🔹 Properties of Metal Chlorides

Ionic (electrovalent) compounds – consist of metal cations and chloride anions.
Solid at room temperature.
High melting and boiling points → non-volatile.
Electrolytes → conduct electricity when molten or in aqueous solution.
Usually soluble in water.

🔹 Examples of Reactions

1. Sodium + Chlorine


2Na (s) + Cl₂ (g) → 2NaCl (s)

2. Calcium + Chlorine


Ca (s) + Cl₂ (g) → CaCl₂ (s)

3. Magnesium + Chlorine


Mg (s) + Cl₂ (g) → MgCl₂ (s)

4. Aluminium + Chlorine


2Al (s) + 3Cl₂ (g) → 2AlCl₃ (s)

5. Zinc + Chlorine


Zn (s) + Cl₂ (g) → ZnCl₂ (s)

6. Iron + Chlorine


2Fe (s) + 3Cl₂ (g) → 2FeCl₃ (s)

7. Copper + Chlorine


Cu (s) + Cl₂ (g) → CuCl₂ (s)

🔹 Key Points to Remember

All metal chlorides are ionic compounds (electrovalent).
Their bonding is due to electron transfer (not sharing, unlike covalent compounds).
Reactivity depends on the metal:
- Highly reactive metals (e.g., Na, Ca, Mg) → react vigorously.
- Less reactive metals (e.g., Cu, Fe) → require heating.

✅ These notes cover definition, properties, reactions, and examples — perfect for exams.

⚡ Reaction of Metals with Hydrogen

🔹 General Rule

Most metals do not react with hydrogen.
Reason:
- Metals form compounds by losing electrons.
- Hydrogen also tends to lose an electron (H⁺) or share electrons.
- Since both want to lose electrons, hydrogen generally does not combine with metals.

👉 Only a few very reactive metals can force hydrogen atoms to accept electrons → forming metal hydrides.

🔹 What are Metal Hydrides?

Salt-like solid compounds formed by some reactive metals with hydrogen.
Ionic in nature.
Contain the hydride ion (H⁻) → hydrogen acts as an anion.
General formula: MH (for alkali metals), MH₂ (for alkaline earth metals).

🔹 Examples of Reactions

1. Sodium + Hydrogen


2Na (s) + H₂ (g) → 2NaH (s)

Ionic compound: Na⁺ + H⁻.

2. Potassium + Hydrogen


2K (s) + H₂ (g) → 2KH (s)

3. Calcium + Hydrogen


Ca (s) + H₂ (g) → CaH₂ (s)

Ionic compound: Ca²⁺ + 2H⁻.

4. Magnesium + Hydrogen


Mg (s) + H₂ (g) → MgH₂ (s)

🔹 Metals That Do Not React with Hydrogen

Less reactive metals (like zinc, iron, copper, silver, gold) do not form hydrides with hydrogen.

🔹 Properties of Metal Hydrides

Ionic compounds → consist of metal cations and hydride anions (H⁻).
Solid at room temperature.
React with water to give hydrogen gas.
Example:


   NaH + H₂O → NaOH + H₂ ↑

🔹 Key Points to Remember

Only very reactive metals (Na, K, Ca, Mg) form hydrides.
In metal hydrides, hydrogen carries a negative charge (H⁻).
Less reactive metals do not combine with hydrogen.

✅ These notes include definition, examples, properties, and key points for exams.

🌟 Chemical Properties of Non-Metals

1️⃣ Reaction of Non-Metals with Oxygen

Non-metals react with oxygen to form non-metal oxides.
These oxides can be:
- Acidic oxides (dissolve in water to form acids, turn blue litmus red).
- Neutral oxides (neither acidic nor basic, no effect on litmus).
Non-metal oxides are covalent compounds (formed by sharing of electrons).

🔹 Examples of Acidic Oxides

(i) Carbon (C)


C (s) + O₂ (g) → CO₂ (g)

Dissolves in water to form carbonic acid (H₂CO₃).
Turns blue litmus red.

(ii) Sulphur (S)


S (s) + O₂ (g) → SO₂ (g)

Dissolves in water to form sulphurous acid (H₂SO₃).
Turns blue litmus red.

(iii) Phosphorus (P)


4P (s) + 5O₂ (g) → 2P₂O₅ (s)

Forms phosphoric acid (H₃PO₄) in water.

👉 Acidic oxides of non-metals are also called acid anhydrides.

🔹 Examples of Neutral Oxides

(i) Carbon monoxide (CO)


2C (s) + O₂ (g) → 2CO (g)  \quad (\text{limited air})

(ii) Water (H₂O)


2H₂ (g) + O₂ (g) → 2H₂O (l)

(iii) Nitrogen monoxide (NO) and Dinitrogen monoxide (N₂O)

Both are neutral oxides → no effect on litmus.

2️⃣ Key Points about Non-Metal Oxides

Acidic Oxides: e.g. CO₂, SO₂, P₂O₅ → dissolve in water to form acids.
Neutral Oxides: e.g. CO, H₂O, NO, N₂O → neither acidic nor basic.
Non-metal oxides are covalent in nature (not ionic).
They do not contain O²⁻ ions.

📝 Quick Summary Table

Non-Metal	Reaction with O₂	Oxide Formed	Nature of Oxide
Carbon (C)	C + O₂ → CO₂	CO₂	Acidic
Sulphur (S)	S + O₂ → SO₂	SO₂	Acidic
Phosphorus (P)	4P + 5O₂ → 2P₂O₅	P₂O₅	Acidic
Carbon (C, limited O₂)	2C + O₂ → 2CO	CO	Neutral
Hydrogen (H₂)	2H₂ + O₂ → 2H₂O	H₂O	Neutral
Nitrogen (N₂)	N₂ + O₂ → 2NO	NO	Neutral

✅ These notes cover definitions, examples, equations, explanations, and quick summary for exams.

🌟 Reactions of Non-Metals

1️⃣ Reaction of Non-Metals with Water 💧

Non-metals do not react with water or steam to evolve hydrogen gas.
Reason:
- To liberate H₂ from water, electrons must be given to H⁺ ions.
- Non-metals are electron acceptors (not donors), so they cannot reduce H⁺ into H₂.
Result: No hydrogen gas is evolved when non-metals are put in water.

2️⃣ Reaction of Non-Metals with Dilute Acids ⚗️

Non-metals do not react with dilute acids (like HCl, H₂SO₄).
Reason:
- Acids contain H⁺ ions. To produce H₂ gas, non-metals must supply electrons.
- But non-metals cannot give electrons (they only accept electrons).
Example:
- Carbon (C), Sulphur (S), Phosphorus (P) do not react with HCl or H₂SO₄.

👉 Conclusion: Non-metals cannot displace hydrogen from acids.

3️⃣ Reaction of Non-Metals with Salt Solutions 🧪

A more reactive non-metal can displace a less reactive non-metal from its salt solution.

Example:


2NaBr (aq) + Cl₂ (g) → 2NaCl (aq) + Br₂ (aq)

👉 This follows the reactivity series of non-metals (F₂ > Cl₂ > Br₂ > I₂).

4️⃣ Reaction of Non-Metals with Chlorine 🟢

Non-metals react with chlorine to form covalent chlorides.
Properties:
- Usually liquids or gases.
- Do not conduct electricity (non-electrolytes).

Examples:

Hydrogen + Chlorine → Hydrogen chloride


H₂ + Cl₂ → 2HCl (g)

Phosphorus + Chlorine → Phosphorus trichloride


P₄ + 6Cl₂ → 4PCl₃ (l)

Carbon + Chlorine → Carbon tetrachloride


C + 2Cl₂ → CCl₄

👉 Non-metals form covalent chlorides because they cannot give electrons to chlorine.

5️⃣ Reaction of Non-Metals with Hydrogen 🔺

Non-metals combine with hydrogen to form covalent hydrides.
Properties:
- Formed by sharing of electrons (covalent bond).
- Generally liquids or gases.
- Do not conduct electricity (no ions).
- Stable compounds.

Examples:

Sulphur + Hydrogen → Hydrogen sulphide


H₂ + S → H₂S (g)

Nitrogen + Hydrogen → Ammonia (with Fe catalyst)


N₂ + 3H₂ \xrightarrow{Fe} 2NH₃ (g)

Oxygen + Hydrogen → Water


2H₂ + O₂ → 2H₂O

Carbon + Hydrogen → Methane


C + 2H₂ → CH₄

👉 Non-metal hydrides are covalent, non-conductors, and stable.

📝 Quick Comparison: Metals vs. Non-Metals

Property	Metals	Non-Metals
Reaction with Water	Displace H₂ from water/steam	No reaction
Reaction with Dilute Acids	Displace H₂	No reaction
Reaction with Chlorine	Form ionic chlorides (electrolytes)	Form covalent chlorides (non-electrolytes)
Reaction with Hydrogen	Rare, form ionic hydrides	Common, form covalent hydrides
Nature of Oxides	Basic oxides	Acidic or neutral oxides

✅ These notes now cover water, acids, salt solutions, chlorine, hydrogen, with reactions and reasoning.

Uses of Metals

✅ Metals are widely used in our daily life due to their properties such as malleability, ductility, strength, conductivity, and resistance to corrosion.

1. Electrical Uses

Copper & Aluminium → used to make electrical wires because they are good conductors of electricity and have low resistance.

2. Household and Industrial Uses

Iron, Copper & Aluminium → used to make utensils, factory equipment, machines, and tools.
Steel bins → galvanised with zinc to prevent rusting.
Chromium → used for electroplating to make objects shiny and rust-resistant.

3. Catalysts and Chemical Uses

Iron → used as a catalyst in the manufacture of ammonia (Haber’s process).
Zinc → used for galvanising iron to protect it from rusting.

4. Alloy & Special Uses

Chromium & Nickel → used in making stainless steel (strong, shiny, rust-resistant).
Zirconium → used in bullet-proof alloy steels.
Sodium, Titanium & Zirconium → used in atomic energy & space research projects.

5. Packaging and Decoration

Aluminium foils → used for packaging medicines, food, and cigarettes.
Gold & Silver → used in jewellery, ornaments, and in thin foils for decorating sweets.

6. Other Important Uses

Mercury → used in thermometers and other instruments.
Lead → used in car batteries.

Uses of Non-Metals

✅ Non-metals are important in fuel, industry, medicine, food preservation, and explosives.

1. Uses of Hydrogen

Used in hydrogenation of vegetable oils to make ghee (vanaspati).
Used in making ammonia (fertilisers).
Liquid hydrogen → used as rocket fuel.

2. Uses of Carbon

Graphite → used to make electrodes in batteries and electrolytic cells.

3. Uses of Nitrogen

Used in making ammonia, nitric acid, and fertilisers.
Used to preserve food materials (due to inert nature).
Compounds like TNT (Tri Nitro Toluene) and nitroglycerine are used as explosives.

4. Uses of Sulphur

Used to manufacture sulphuric acid (an important industrial chemical).
Used as a fungicide and in making gunpowder.
Used in the vulcanisation of rubber (to make rubber strong and elastic).

Quick Summary Table

Metals	Uses
Copper, Aluminium	Wires, utensils, packaging
Iron	Machines, catalyst in Haber’s process
Zinc	Galvanisation (rust protection)
Chromium, Nickel	Stainless steel, electroplating
Silver, Gold	Jewellery, decoration
Mercury	Thermometers
Lead	Car batteries

Non-Metals	Uses
Hydrogen	Rocket fuel, ammonia, hydrogenation of oils
Carbon (Graphite)	Electrodes, dry cells
Nitrogen	Fertilisers, explosives, food preservation
Sulphur	Sulphuric acid, fungicide, gunpowder, rubber vulcanisation

✨ Tip for Exams:

Always connect the property of the metal/non-metal with its use.
Example: Copper is used in wires because it is a good conductor of electricity.