Introduction

Everything around us—from the spoon we use to eat food to the oxygen we breathe—consists of elements. These elements are broadly classified into metals and non-metals based on their physical and chemical properties. Metals are known for their strength, lustre, and ability to conduct heat and electricity, whereas non-metals generally have opposite characteristics.

Metals play a vital role in construction, transportation, electrical wiring, machinery, and household appliances. Non-metals are equally important because they are essential for life processes, agriculture, medicine, and various industries. Understanding the differences between metals and non-metals helps us choose the right materials for different applications.

In this chapter, you will learn about the physical and chemical properties of metals and non-metals, the reactivity series, ionic compounds, extraction of metals, corrosion, and methods of preventing corrosion.

Learning Objectives

After studying this chapter, you will be able to:

• Define metals and non-metals.
• Differentiate between metals and non-metals based on their properties.
• Explain the physical and chemical properties of metals.
• Describe the chemical properties of non-metals.
• Understand the concept of the reactivity series.
• Explain the formation and properties of ionic compounds.
• Learn how metals are extracted from their ores.
• Understand corrosion and methods to prevent it.

What are Metals?

Metals are elements that generally possess properties such as lustre, hardness, malleability, ductility, and high electrical and thermal conductivity.

Most metals are found in the left and centre of the periodic table.

Examples include:

• Iron (Fe)
• Copper (Cu)
• Aluminium (Al)
• Gold (Au)
• Silver (Ag)
• Zinc (Zn)
• Sodium (Na)
• Magnesium (Mg)

Metals are widely used because of their excellent mechanical and physical properties.

Everyday Uses of Metals

Metals are an important part of our daily lives.

What are Non-Metals?

Non-metals are elements that generally lack the typical properties of metals. They are usually poor conductors of heat and electricity and are neither malleable nor ductile.

Most non-metals are found on the right side of the periodic table.

Examples include:

• Oxygen (O)
• Nitrogen (N)
• Carbon (C)
• Sulphur (S)
• Chlorine (Cl)
• Phosphorus (P)
• Bromine (Br)
• Iodine (I)

Non-metals are essential for life and many industrial processes.

Everyday Uses of Non-Metals

Physical Properties of Metals

Metals exhibit several characteristic physical properties that distinguish them from non-metals.

1. Lustre

Metals have a shiny surface known as metallic lustre.

Examples:

• Gold jewellery shines brightly.
• Silver utensils have a glossy appearance.
• Aluminium foil reflects light.

Why are metals lustrous?

The free electrons present in metals reflect light, giving them a shiny appearance.

2. Hardness

Most metals are hard and difficult to cut.

Examples:

• Iron
• Copper
• Nickel

However, some metals such as sodium and potassium are soft and can be cut with a knife.

3. Malleability

Definition

The property by which metals can be beaten into thin sheets is called malleability.

Examples:

• Aluminium sheets
• Gold leaf
• Silver foil

Applications

• Aluminium foil used for food packaging
• Gold leaf used in decoration
• Metal sheets used in vehicle manufacturing

Gold is one of the most malleable metals.

4. Ductility

Definition

The property by which metals can be drawn into thin wires is called ductility.

Examples:

• Copper wires
• Aluminium wires
• Gold wires

Applications

• Electrical wiring
• Telephone cables
• Motor windings

Gold is also one of the most ductile metals.

5. Conductivity

Metals conduct both heat and electricity efficiently because they contain free-moving electrons.

Good Conductors

• Silver (best conductor of electricity)
• Copper
• Aluminium

Applications

• Electrical wires
• Electric motors
• Cooking utensils
• Heat exchangers

6. Sonority

Definition

The property of producing a ringing sound when struck is called sonority.

Examples:

• School bells
• Temple bells
• Brass musical instruments

This property is unique to metals.

7. High Melting and Boiling Points

Most metals have high melting and boiling points.

Examples:

• Iron melts at approximately 1538°C.
• Copper melts at approximately 1085°C.

This makes metals suitable for machinery and construction.

8. High Density

Most metals are heavy due to their high density.

Examples:

• Iron
• Copper
• Lead

9. Strength

Metals are generally strong and can withstand heavy loads.

Applications:

• Bridges
• Buildings
• Railway tracks
• Machines

Physical Properties of Non-Metals

Non-metals generally show properties opposite to those of metals.

1. Non-Lustrous

Most non-metals have a dull appearance.

Examples:

• Sulphur
• Phosphorus
• Carbon (coal)

Exception: Iodine has a shiny surface despite being a non-metal.

2. Brittle

Solid non-metals break easily when hammered.

They cannot be beaten into sheets.

Examples:

• Sulphur
• Phosphorus

3. Non-Malleable

Non-metals cannot be hammered into thin sheets because they are brittle.

4. Non-Ductile

They cannot be drawn into wires.

5. Poor Conductors

Most non-metals are poor conductors of heat and electricity.

Examples:

• Sulphur
• Phosphorus
• Oxygen

Exception

Graphite, an allotrope of carbon, conducts electricity and is used in electrodes and batteries.

6. Non-Sonorous

Non-metals do not produce a ringing sound when struck.

7. Low Density

Many non-metals are lighter than metals.

Examples:

• Carbon
• Sulphur
• Phosphorus

8. Low Melting and Boiling Points

Many non-metals melt and boil at relatively low temperatures.

Examples:

• Sulphur
• Phosphorus

Comparison Between Metals and Non-Metals

Exceptions to General Properties

Not all elements follow the general rules. These exceptions are frequently asked in CBSE examinations.

Activity: Testing Malleability

Materials Required

• Copper wire
• Aluminium strip
• Sulphur piece
• Small hammer

Procedure

1. Place each sample on a hard surface.
2. Gently strike it with the hammer.
3. Observe the changes.

Observation

• Copper and aluminium flatten into sheets.
• Sulphur breaks into small pieces.

Conclusion

Metals are malleable, whereas non-metals are brittle.

Did You Know?

🥇 Gold is so malleable that one gram of gold can be beaten into a sheet covering nearly one square metre. This unique property makes it ideal for jewellery and decorative applications.

Exam Tip

Do not memorize all properties separately. Instead, remember the key pattern:

• Metals: Lustrous, hard, malleable, ductile, sonorous, good conductors.
• Non-metals: Generally dull, brittle, non-sonorous, poor conductors.

Knowing the exceptions—such as graphite, iodine, bromine, mercury, sodium, and potassium—is equally important for scoring well in CBSE board examinations.

Quick Revision

• Metals are generally lustrous, hard, malleable, ductile, sonorous, and good conductors of heat and electricity.
• Non-metals are usually dull, brittle, non-malleable, non-ductile, non-sonorous, and poor conductors.
• Graphite is a non-metal that conducts electricity.
• Iodine is a lustrous non-metal.
• Mercury is the only common metal that is liquid at room temperature.
• Bromine is the only common non-metal that is liquid at room temperature.
• Gold is one of the most malleable and ductile metals.

Chemical Properties of Metals

The physical properties of metals help us identify them, but their chemical properties explain how they react with other substances. Metals generally lose electrons during chemical reactions and form positive ions (cations). This is why metals are called electropositive elements.

The important chemical properties of metals include their reactions with:

• Oxygen
• Water
• Dilute acids
• Salt solutions
• Non-metals

These reactions help us understand the reactivity of different metals.

1. Reaction of Metals with Oxygen

Most metals react with oxygen to form metal oxides.

General Equation

Metal + Oxygen → Metal Oxide

Most metal oxides are basic in nature.

Example 1: Magnesium

When magnesium ribbon is heated in air, it burns with a dazzling white flame and forms a white powder of magnesium oxide.

Balanced Equation

$$2Mg + O_2 \rightarrow 2MgO$$2Mg+O2​→2MgO

Observation

• Bright white flame
• White ash is formed
• Heat and light are released

Example 2: Copper

Copper does not burn easily but on heating, it reacts with oxygen to form black copper(II) oxide.

Balanced Equation

$$2Cu + O_2 \rightarrow 2CuO$$2Cu+O2​→2CuO

Observation

• Brown copper turns black.

Example 3: Aluminium

Aluminium reacts with oxygen to form a thin layer of aluminium oxide.

Balanced Equation

$$4Al + 3O_2 \rightarrow 2Al_2O_3$$4Al+3O2​→2Al2​O3​

The oxide layer protects aluminium from further corrosion.

Nature of Metal Oxides

Most metal oxides are basic oxides because they react with acids to form salt and water.

Example

$$CuO + 2HCl \rightarrow CuCl_2 + H_2O$$CuO+2HCl→CuCl2​+H2​O

Copper oxide reacts with hydrochloric acid to form copper chloride and water.

Amphoteric Oxides

Some metal oxides react with both acids and bases.

Such oxides are called amphoteric oxides.

Examples

• Aluminium oxide (Al₂O₃)
• Zinc oxide (ZnO)

Aluminium Oxide with Acid

$$Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$$Al2​O3​+6HCl→2AlCl3​+3H2​O

Aluminium Oxide with Base

$$Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O$$Al2​O3​+2NaOH→2NaAlO2​+H2​O

This dual behaviour makes aluminium oxide amphoteric.

2. Reaction of Metals with Water

Different metals react with water at different rates.

Some react violently, while others react very slowly or not at all.

A. Sodium and Potassium

These metals react very vigorously with cold water.

Example

$$2Na + 2H_2O \rightarrow 2NaOH + H_2 + \text{Heat}$$2Na+2H2​O→2NaOH+H2​+Heat

Observation

• Hydrogen gas evolves rapidly.
• Large amount of heat is released.
• Sodium floats and moves rapidly on water.
• The reaction may catch fire.

Because of their high reactivity, sodium and potassium are stored in kerosene oil.

B. Calcium

Calcium reacts less vigorously than sodium.

Equation

$$Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2$$Ca+2H2​O→Ca(OH)2​+H2​

Observation

• Hydrogen gas is released.
• Calcium floats due to bubbles sticking to its surface.

C. Magnesium

Magnesium reacts very slowly with cold water but reacts rapidly with hot water and steam.

With Steam

$$Mg + H_2O(\text{Steam}) \rightarrow MgO + H_2$$Mg+H2​O(Steam)→MgO+H2​

D. Iron

Iron does not react with cold or hot water but reacts with steam.

Equation

$$3Fe + 4H_2O(\text{Steam}) \rightarrow Fe_3O_4 + 4H_2$$3Fe+4H2​O(Steam)→Fe3​O4​+4H2​

E. Copper, Silver and Gold

These metals do not react with water because they are less reactive.

Reaction of Metals with Water (Summary Table)

3. Reaction of Metals with Dilute Acids

Most metals react with dilute acids to produce salt and hydrogen gas.

General Equation

Metal + Dilute Acid → Salt + Hydrogen Gas

Example 1

$$Zn + 2HCl \rightarrow ZnCl_2 + H_2$$Zn+2HCl→ZnCl2​+H2​

Hydrogen gas is released.

Example 2

$$Mg + H_2SO_4 \rightarrow MgSO_4 + H_2$$Mg+H2​SO4​→MgSO4​+H2​

Observation

• Effervescence (bubbles) is seen.
• Hydrogen gas burns with a ‘pop’ sound when tested using a burning splint.

Why Do Some Metals Not React?

Metals like:

• Copper
• Silver
• Gold

are less reactive than hydrogen and therefore do not displace hydrogen from dilute acids.

4. Reaction of Metals with Salt Solutions

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

This is called a displacement reaction.

Example 1

$$Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$$Zn+CuSO4​→ZnSO4​+Cu

Observation

• Blue copper sulphate solution becomes colourless.
• Copper is deposited.

Example 2

$$Fe + CuSO_4 \rightarrow FeSO_4 + Cu$$Fe+CuSO4​→FeSO4​+Cu

Observation

• Blue solution turns green.
• Copper is deposited on the iron nail.

Example 3

Copper cannot displace zinc from zinc sulphate solution.$$Cu + ZnSO_4 \rightarrow \text{No Reaction}$$Cu+ZnSO4​→No Reaction

Because copper is less reactive than zinc.

Reactivity Series

The reactivity series is a list of metals arranged in decreasing order of their reactivity.

Reactivity Series

Importance of Reactivity Series

The reactivity series helps us predict:

• Which metal reacts faster.
• Which metal can displace another.
• Which metals react with water.
• Which metals react with acids.
• Which extraction method should be used.
• Which metals are more resistant to corrosion.

Memory Trick

A commonly used mnemonic is:

“Please Stop Calling Me A Zebra Instead Try Learning How Copper Saves Gold.”

It represents:

• Potassium
• Sodium
• Calcium
• Magnesium
• Aluminium
• Zinc
• Iron
• Tin (if included in extended series)
• Lead
• Hydrogen
• Copper
• Silver
• Gold

(Note: Different textbooks may use slightly different mnemonics. Always follow your CBSE syllabus.)

Activity: Displacement Reaction

Materials

• Iron nail
• Copper sulphate solution
• Beaker

Procedure

1. Clean the iron nail with sandpaper.
2. Place it in copper sulphate solution.
3. Leave it for 20–30 minutes.

Observation

• The blue colour of the solution fades.
• A reddish-brown coating of copper appears on the iron nail.

Conclusion

Iron is more reactive than copper and displaces it from copper sulphate solution.

Everyday Applications

Exam Tip

Remember these key trends:

• Metals above hydrogen in the reactivity series react with dilute acids to produce hydrogen gas.
• Metals below hydrogen (such as copper, silver, and gold) do not react with dilute acids.
• A more reactive metal always displaces a less reactive metal from its salt solution.

Did You Know?

🚀 Aluminium appears highly reactive, yet it does not corrode easily. This is because a thin, tough layer of aluminium oxide forms on its surface, preventing further reaction with air and moisture. This property makes aluminium ideal for aircraft bodies, window frames, and kitchen utensils.

Quick Revision

• Metals react with oxygen to form metal oxides, which are usually basic.
• Aluminium oxide and zinc oxide are amphoteric oxides.
• Highly reactive metals react vigorously with water.
• Most metals react with dilute acids to release hydrogen gas.
• A more reactive metal displaces a less reactive metal from its salt solution.
• The reactivity series arranges metals from most reactive to least reactive and helps predict their chemical behaviour.

Chemical Properties of Non-Metals

Unlike metals, non-metals generally gain electrons during chemical reactions and form negative ions (anions). Therefore, non-metals are called electronegative elements.

Non-metals react with oxygen, hydrogen, and metals to form different compounds that are widely used in daily life and industries.

The important chemical properties of non-metals include their reactions with:

• Oxygen
• Hydrogen
• Metals

1. Reaction of Non-Metals with Oxygen

Most non-metals react with oxygen to form non-metal oxides.

General Equation

Non-Metal + Oxygen → Non-Metal Oxide

Most non-metal oxides are acidic in nature.

Example 1: Carbon

$$C + O_2 \rightarrow CO_2$$C+O2​→CO2​

Carbon burns in oxygen to form carbon dioxide.

Example 2: Sulphur

$$S + O_2 \rightarrow SO_2$$S+O2​→SO2​

Sulphur burns with a blue flame and forms sulphur dioxide.

Observation

• Blue flame
• Pungent smell
• Colourless gas produced

Example 3: Phosphorus

$$4P + 5O_2 \rightarrow 2P_2O_5$$4P+5O2​→2P2​O5​

Phosphorus reacts vigorously with oxygen.

Nature of Non-Metal Oxides

Most non-metal oxides are acidic oxides.

They react with bases to form salt and water.

Example

$$CO_2 + Ca(OH)_2 \rightarrow CaCO_3 + H_2O$$CO2​+Ca(OH)2​→CaCO3​+H2​O

Carbon dioxide reacts with lime water to form calcium carbonate.

Neutral Oxides

Some oxides are neither acidic nor basic.

These are called neutral oxides.

Examples:

• Carbon monoxide (CO)
• Nitric oxide (NO)

2. Reaction of Non-Metals with Hydrogen

Many non-metals combine with hydrogen to form covalent compounds.

Examples

Hydrogen + Chlorine$$H_2 + Cl_2 \rightarrow 2HCl$$H2​+Cl2​→2HCl

Hydrogen + Sulphur$$H_2 + S \rightarrow H_2S$$H2​+S→H2​S

These compounds are important in industries and laboratories.

3. Reaction of Non-Metals with Metals

Non-metals react with metals to form ionic compounds.

Example

$$2Na + Cl_2 \rightarrow 2NaCl$$2Na+Cl2​→2NaCl

Sodium reacts with chlorine to form sodium chloride (common salt).

What are Ionic Compounds?

Ionic compounds are compounds formed by the transfer of electrons from a metal atom to a non-metal atom.

They are also called electrovalent compounds.

Ionic Bond

Definition

An ionic bond is the force of attraction between positively charged ions (cations) and negatively charged ions (anions).

Formation of Sodium Chloride (NaCl)

This is the most important example in the CBSE syllabus.

Step 1: Sodium Atom

Electronic configuration: 2,8,1

Sodium has one electron in its outermost shell.

It loses one electron.$$Na \rightarrow Na^+ + e^-$$Na→Na++e−

A positive sodium ion is formed.

Step 2: Chlorine Atom

Electronic configuration: 2,8,7

Chlorine requires one electron to complete its octet.

It gains one electron.$$Cl + e^- \rightarrow Cl^-$$Cl+e−→Cl−

A negative chloride ion is formed.

Step 3: Formation of Ionic Bond

The opposite charges attract each other.$$Na^+ + Cl^- \rightarrow NaCl$$Na++Cl−→NaCl

This electrostatic attraction forms an ionic bond.

Why Do Atoms Form Ionic Compounds?

Atoms combine to achieve a stable electronic configuration similar to that of noble gases.

This is called the octet rule.

Atoms become stable when they have:

• 8 electrons in the outermost shell
• (Except hydrogen and helium)

Characteristics of Ionic Compounds

1. Solid and Hard

Ionic compounds are usually hard crystalline solids.

Examples:

• Sodium chloride
• Potassium chloride

2. High Melting and Boiling Points

Strong electrostatic forces hold ions together.

Large amounts of heat are required to break these forces.

Examples:

• NaCl melts at about 801°C.

3. Soluble in Water

Most ionic compounds dissolve easily in water.

Examples:

• NaCl
• KCl

However, many ionic compounds are insoluble in organic solvents like petrol and kerosene.

4. Conduct Electricity in Molten or Aqueous State

Solid ionic compounds do not conduct electricity because ions cannot move freely.

However,

• Molten ionic compounds conduct electricity.
• Aqueous solutions also conduct electricity.

This is because ions become free to move.

5. Brittle

When struck with a hammer, ionic crystals break into pieces.

They are hard but brittle.

Properties of Ionic Compounds (Summary Table)

Metals vs Non-Metals (Chemical Behaviour)

Activity: Electrical Conductivity of Ionic Compounds

Materials

• Battery
• Bulb
• Wires
• Salt solution
• Solid salt

Procedure

1. Connect the bulb to the battery.
2. Test solid sodium chloride.
3. Test sodium chloride solution.

Observation

• Solid NaCl does not light the bulb.
• Salt solution lights the bulb.

Conclusion

Ions conduct electricity only when they are free to move.

Uses of Ionic Compounds

Sodium Chloride (NaCl)

• Food seasoning
• Food preservation
• Chemical industry

Calcium Chloride (CaCl₂)

• Drying agent
• De-icing roads

Magnesium Oxide (MgO)

• Refractory bricks
• Furnace lining

Potassium Chloride (KCl)

• Fertilisers

Everyday Examples

Exam Tip

Remember this concept clearly:

• Metals lose electrons → Cations (positive ions).
• Non-metals gain electrons → Anions (negative ions).
• Transfer of electrons forms an ionic bond.

This is one of the most frequently asked concepts in CBSE board examinations.

Did You Know?

🧂 Common salt (NaCl) is one of the most important ionic compounds. Besides making food tasty, it is used in the manufacture of chlorine gas, caustic soda (NaOH), hydrochloric acid (HCl), soaps, detergents, and many other industrial chemicals.

Occurrence of Metals

Metals are found naturally in the Earth’s crust. Some metals occur in the free (native) state, while others are found in the combined state as compounds.

The form in which a metal occurs depends on its reactivity.

• Highly reactive metals (e.g., sodium, potassium) are always found in the combined state.
• Moderately reactive metals (e.g., iron, zinc) are usually found as oxides, sulphides, or carbonates.
• Less reactive metals (e.g., gold, silver, platinum) may occur in the free state.

Minerals

A mineral is a naturally occurring substance in the Earth’s crust that contains metals or their compounds.

Examples

• Bauxite
• Hematite
• Galena
• Cinnabar
• Chalcopyrite

Ores

An ore is a mineral from which a metal can be extracted economically and conveniently.

Every ore is a mineral, but every mineral is not an ore.

Common Ores of Important Metals

Metallurgy

The process of extracting pure metal from its ore is called metallurgy.

The main steps in metallurgy are:

1. Mining
2. Crushing and Grinding
3. Concentration (Enrichment) of Ore
4. Extraction of Metal
5. Refining (Purification)

Step 1: Mining

Mining is the process of removing ores from the Earth’s crust.

Types of Mining

A. Open-Cast Mining

• Used when ores are close to the surface.
• Less expensive.
• Safer than deep mining.

B. Underground Mining

• Used for deep deposits.
• More expensive.
• Requires tunnels and shafts.

Step 2: Crushing and Grinding

Large rocks are crushed into smaller pieces.

This increases the surface area and makes extraction easier.

Step 3: Concentration (Enrichment) of Ore

Ores contain unwanted materials such as sand, clay, and stones.

These unwanted substances are called gangue or impurities.

Concentration removes these impurities.

Methods of Concentration

A. Hydraulic Washing

Uses water to separate heavier ore particles from lighter impurities.

Suitable for heavy oxide ores.

B. Froth Flotation Process

Used mainly for sulphide ores.

Air is passed through the powdered ore mixed with water and oil.

Sulphide particles attach to the froth and rise to the surface.

C. Magnetic Separation

Used when the ore is magnetic but impurities are not, or vice versa.

Step 4: Extraction of Metals

Different extraction methods are used depending on the reactivity of the metal.

A. Extraction of Highly Reactive Metals

Examples:

• Sodium
• Potassium
• Calcium
• Magnesium
• Aluminium

These metals cannot be extracted using carbon because they are more reactive than carbon.

They are extracted by electrolysis.

Electrolysis

Electric current is passed through the molten compound.

The metal is deposited at the cathode.

Example

Extraction of sodium from molten sodium chloride.$$2NaCl \xrightarrow{\text{Electricity}} 2Na + Cl_2$$2NaClElectricity​2Na+Cl2​

B. Extraction of Moderately Reactive Metals

Examples:

• Iron
• Zinc
• Lead

These metals are extracted by reduction with carbon (coke).

Example

Extraction of iron$$Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$$Fe2​O3​+3CO→2Fe+3CO2​

Carbon monoxide removes oxygen from iron oxide.

C. Extraction of Least Reactive Metals

Examples:

• Mercury
• Silver
• Gold

These metals occur in native form or require simple heating.

Example

Extraction of mercury$$2HgO \xrightarrow{\Delta} 2Hg + O_2$$2HgOΔ​2Hg+O2​

Refining of Metals

The metal obtained after extraction is often impure.

The process of removing impurities is called refining.

Electrolytic Refining

This is one of the most common purification methods.

Components

• Impure metal → Anode
• Pure metal sheet → Cathode
• Metal salt solution → Electrolyte

During electrolysis:

• Pure metal deposits on the cathode.
• Impurities settle as anode mud.

This method is commonly used for:

• Copper
• Silver
• Gold

Corrosion

Definition

Corrosion is the gradual destruction of metals due to chemical reactions with air, moisture, acids, or other environmental substances.

Rusting of Iron

Rusting is the most familiar example of corrosion.

Chemical Equation

$$4Fe + 3O_2 + xH_2O \rightarrow 2Fe_2O_3 \cdot xH_2O$$4Fe+3O2​+xH2​O→2Fe2​O3​⋅xH2​O

Rust is hydrated iron(III) oxide.

Conditions Necessary for Rusting

Rusting occurs only when both are present:

• Oxygen
• Moisture (Water)

Without either of these, rusting does not take place.

Harmful Effects of Corrosion

Corrosion causes:

• Damage to buildings
• Weakening of bridges
• Leakage in pipelines
• Damage to ships
• Reduced life of vehicles
• Failure of machinery
• Economic loss

Prevention of Corrosion

Several methods are used to protect metals from corrosion.

1. Painting

Paint forms a protective layer that prevents contact with air and moisture.

Examples:

• Gates
• Bridges
• Railings

2. Oiling and Greasing

Oil and grease protect machine parts.

Examples:

• Bicycle chains
• Engine parts
• Industrial machines

3. Galvanisation

Galvanisation is the process of coating iron with a thin layer of zinc.

Advantages:

• Prevents rusting
• Increases life of iron

Applications:

• Water pipes
• Buckets
• Roofing sheets
• Iron poles

4. Electroplating

A thin layer of another metal is deposited over the object.

Examples:

• Gold plating
• Silver plating
• Chromium plating

Benefits:

• Improves appearance
• Prevents corrosion
• Increases durability

5. Alloying

An alloy is a homogeneous mixture of two or more elements, at least one of which is a metal.

Alloys are generally:

• Stronger
• Harder
• More resistant to corrosion
• More durable

Common Alloys

Why are Alloys Preferred?

Alloys are preferred because they:

• Resist corrosion
• Have greater strength
• Last longer
• Have better appearance
• Can be designed for specific purposes

Everyday Applications

Activity: Rusting of Iron

Materials

• Three test tubes
• Iron nails
• Water
• Calcium chloride
• Boiled water
• Oil

Procedure

1. Place an iron nail in each test tube.
2. First tube: Water + Air
3. Second tube: Dry air (using calcium chloride)
4. Third tube: Boiled water covered with oil

Observation

• Rust forms only in the first test tube.

Conclusion

Rusting requires both oxygen and water.

Exam Tip

Remember this order of extraction:

• Highly reactive metals → Electrolysis
• Moderately reactive metals → Reduction using carbon
• Least reactive metals → Heating or found in native state

This classification is frequently tested in CBSE board exams.

Did You Know?

🏗️ The Eiffel Tower in Paris is repainted approximately every seven years to protect its iron structure from corrosion. Regular maintenance helps preserve one of the world’s most famous landmarks.

Quick Revision

• Metals occur in the Earth as minerals and ores.
• Ores are minerals from which metals can be extracted economically.
• Metallurgy involves mining, concentration, extraction, and refining.
• Highly reactive metals are extracted by electrolysis.
• Moderately reactive metals are extracted by reduction with carbon.
• Corrosion is the gradual destruction of metals.
• Rusting is the corrosion of iron in the presence of oxygen and water.
• Painting, galvanisation, electroplating, oiling, and alloying help prevent corrosion.
• Alloys are stronger and more corrosion-resistant than many pure metals.

Important Definitions

1. Metal

A metal is an element that is generally lustrous, malleable, ductile, sonorous, and a good conductor of heat and electricity.

2. Non-Metal

A non-metal is an element that is generally dull, brittle (if solid), non-sonorous, and a poor conductor of heat and electricity.

3. Mineral

A mineral is a naturally occurring substance in the Earth’s crust containing metals or their compounds.

4. Ore

An ore is a mineral from which a metal can be extracted economically and conveniently.

5. Metallurgy

The process of extracting pure metals from their ores is called metallurgy.

6. Gangue

The unwanted earthy impurities such as sand, clay, and rocks present in an ore are called gangue.

7. Concentration of Ore

The process of removing gangue from an ore is called concentration (or enrichment) of ore.

8. Electrolysis

The process of extracting or purifying metals by passing electricity through a molten compound or solution is called electrolysis.

9. Corrosion

The gradual destruction of metals due to the action of air, moisture, or chemicals is called corrosion.

10. Rusting

The corrosion of iron in the presence of oxygen and moisture, forming hydrated iron(III) oxide, is called rusting.

11. Galvanisation

The process of coating iron or steel with a layer of zinc to prevent rusting is called galvanisation.

12. Electroplating

The process of depositing a thin layer of one metal over another using electricity is called electroplating.

13. Alloy

An alloy is a homogeneous mixture of two or more elements, at least one of which is a metal.

14. Ionic Bond

An ionic bond is the electrostatic force of attraction between positively charged ions and negatively charged ions.

15. Ionic Compound

An ionic compound is a compound formed by the transfer of electrons from a metal to a non-metal.

NCERT Keywords

• Metals
• Non-metals
• Lustre
• Malleability
• Ductility
• Sonority
• Conductivity
• Reactivity Series
• Metal Oxides
• Amphoteric Oxides
• Ionic Bond
• Ionic Compound
• Cation
• Anion
• Minerals
• Ores
• Gangue
• Metallurgy
• Concentration of Ore
• Electrolysis
• Reduction
• Corrosion
• Rusting
• Galvanisation
• Electroplating
• Alloy

Important Chemical Equations

Magnesium Burns in Air

$$2Mg + O_2 \rightarrow 2MgO$$2Mg+O2​→2MgO

Copper Reacts with Oxygen

$$2Cu + O_2 \rightarrow 2CuO$$2Cu+O2​→2CuO

Sodium Reacts with Water

$$2Na + 2H_2O \rightarrow 2NaOH + H_2$$2Na+2H2​O→2NaOH+H2​

Calcium Reacts with Water

$$Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2$$Ca+2H2​O→Ca(OH)2​+H2​

Zinc with Hydrochloric Acid

$$Zn + 2HCl \rightarrow ZnCl_2 + H_2$$Zn+2HCl→ZnCl2​+H2​

Iron with Copper Sulphate

$$Fe + CuSO_4 \rightarrow FeSO_4 + Cu$$Fe+CuSO4​→FeSO4​+Cu

Sodium Chloride Formation

$$2Na + Cl_2 \rightarrow 2NaCl$$2Na+Cl2​→2NaCl

Extraction of Sodium

$$2NaCl \xrightarrow{\text{Electricity}} 2Na + Cl_2$$2NaClElectricity​2Na+Cl2​

Extraction of Iron

$$Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$$Fe2​O3​+3CO→2Fe+3CO2​

Rusting of Iron

$$4Fe + 3O_2 + xH_2O \rightarrow 2Fe_2O_3 \cdot xH_2O$$4Fe+3O2​+xH2​O→2Fe2​O3​⋅xH2​O

Important Differences

Metals vs Non-Metals

Mineral vs Ore

Corrosion vs Rusting

Common Exceptions

These exceptions are very important for CBSE examinations.

One-Page Quick Revision

✔ Metals are generally hard, lustrous, malleable, ductile, sonorous, and good conductors.

✔ Non-metals are generally dull, brittle, non-sonorous, and poor conductors.

✔ Metals lose electrons to form positive ions.

✔ Non-metals gain electrons to form negative ions.

✔ Metal + Oxygen → Metal Oxide (usually basic)

✔ Non-Metal + Oxygen → Non-Metal Oxide (usually acidic)

✔ Ionic compounds are formed by transfer of electrons.

✔ Highly reactive metals are extracted by electrolysis.

✔ Moderately reactive metals are extracted using carbon.

✔ Less reactive metals may occur in the native state.

✔ Corrosion damages metals.

✔ Galvanisation protects iron from rusting.

✔ Alloys are stronger and more corrosion-resistant than pure metals.

Board Exam Tips

1. Learn the Reactivity Series

Questions based on displacement reactions and extraction methods are common.

2. Practice Chemical Equations

Memorise and balance all important reactions.

3. Remember the Exceptions

Graphite, iodine, bromine, mercury, sodium, potassium, aluminium oxide, and zinc oxide are frequently asked in MCQs.

4. Revise Ionic Bond Formation

The formation of sodium chloride is a favourite board exam question.

5. Study Corrosion Carefully

Understand:

• Causes
• Prevention
• Real-life applications

6. Draw Neat Diagrams

Practice diagrams of:

• Ionic bond formation
• Rusting experiment
• Electrolytic refining
• Reactivity series chart

Common Mistakes Students Make

❌ Confusing ores with minerals.

❌ Thinking all metal oxides are basic (remember amphoteric oxides).

❌ Forgetting that solid ionic compounds do not conduct electricity.

❌ Mixing up oxidation and reduction.

❌ Not learning the reactivity series in the correct order.

Frequently Asked Questions (FAQs)

Q1. Why are metals good conductors of electricity?

Metals contain free electrons that move easily and carry electric current.

Q2. Why is sodium stored in kerosene oil?

Sodium reacts vigorously with air and water. Kerosene prevents contact with moisture and oxygen.

Q3. Why is graphite used in electrodes?

Graphite is a non-metal but conducts electricity due to the presence of free electrons.

Q4. What is an alloy?

An alloy is a homogeneous mixture of two or more elements, at least one of which is a metal.

Q5. Why are alloys preferred over pure metals?

Alloys are generally stronger, harder, and more resistant to corrosion.

Q6. What is galvanisation?

It is the process of coating iron or steel with zinc to prevent rusting.

Q7. What is the difference between corrosion and rusting?

Corrosion is the deterioration of any metal, whereas rusting is the corrosion of iron in the presence of oxygen and moisture.

Q8. Which metals are extracted by electrolysis?

Highly reactive metals such as sodium, potassium, calcium, magnesium, and aluminium.

Q9. Why do ionic compounds have high melting points?

Strong electrostatic forces between oppositely charged ions require a large amount of energy to break.

Q10. Which topics are most important for the CBSE Board Exam?

Students should focus on:

• Physical and chemical properties of metals and non-metals
• Reactivity series
• Ionic bond formation
• Extraction of metals
• Corrosion and its prevention
• Important chemical equations
• NCERT activities and in-text questions

Chapter Summary

In this chapter, you explored the fascinating world of metals and non-metals. You learned how their physical and chemical properties differ and why these differences make them suitable for different uses. You studied the reactivity series, which helps predict how metals behave in chemical reactions and guides the methods used to extract them from their ores.

The chapter also introduced the formation of ionic compounds, the concept of ionic bonding, and the industrial process of metallurgy. Finally, you understood the causes of corrosion, methods to prevent it, and the importance of alloys in everyday life. A thorough understanding of these concepts provides a strong foundation for higher chemistry and plays an important role in the CBSE Class 10 Science Board Examination.

💡 Did You Know?

🚗 Modern cars contain dozens of different metals and alloys, including steel, aluminium, copper, magnesium, and titanium. Engineers choose each material carefully to balance strength, weight, safety, durability, and fuel efficiency.

📘 Master CBSE Class 10 Science with the Complete Study Guide

These free notes provide a strong foundation for Chapter 3: Metals and Non-Metals. For complete board exam preparation, explore the CBSE Class 10 Science Master Guide by Science World By Tushar Sir.

🌟 What’s Inside the Complete Book?

• ✅ Complete chapter-wise theory with easy explanations
• ✅ NCERT concepts explained in simple language
• ✅ Mind Maps for quick revision
• ✅ Important definitions and key points
• ✅ Chapter-wise MCQs
• ✅ Assertion & Reason Questions
• ✅ Case-Based Questions
• ✅ Previous Year Board Questions
• ✅ Short & Long Answer Questions
• ✅ Practice Papers
• ✅ Smart Revision Strategies for Board Exams

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