Class 10 Science – Chapter 3 Metals and Non-metals

1. Introduction

Elements are broadly classified into metals and non-metals according to their physical and chemical properties.

Metals are generally useful for making tools, machines, utensils, wires and structures because they are usually strong, malleable, ductile and good conductors. Non-metals have different properties and are important in fuels, fertilizers, medicines, disinfectants and many other applications.

2. Physical Properties of Metals

Important properties

1. Lustre

Metals generally have a shiny surface when freshly prepared or polished.
This property is called metallic lustre.

2. Malleability

The ability of a metal to be beaten into thin sheets is called malleability.
Gold and silver are among the most malleable metals.

3. Ductility

The ability of a metal to be drawn into thin wires is called ductility.
Gold is exceptionally ductile.

4. Hardness

Most metals are hard, although their hardness differs from one metal to another.
Some metals, such as sodium and potassium, are soft enough to be cut with a knife.

5. Conductivity

Metals generally conduct heat and electricity well.
Silver and copper are particularly good conductors of heat.
Copper is widely used in electrical wiring because of its good electrical conductivity.

6. High melting points

Most metals have relatively high melting points.
There are exceptions, such as gallium and caesium, which have comparatively low melting points.

7. Sonority

Metals generally produce a ringing sound when struck.
Such metals are called sonorous.
This is one reason metals are used in bells.

8. Physical state

Almost all metals are solids at room temperature.
Mercury is an important exception because it is liquid at room temperature.

Important exceptions

Physical properties alone cannot perfectly separate metals and non-metals.

Examples:

Iodine is a non-metal but has lustre.
Graphite, a form of carbon, conducts electricity.
Diamond, another form of carbon, is extremely hard.
Sodium, potassium and lithium are unusually soft metals.

3. Physical Properties of Non-metals

Non-metals generally show properties opposite to metals.

They are usually:

Dull rather than lustrous.
Brittle when solid.
Neither malleable nor ductile.
Poor conductors of heat and electricity.
Not sonorous.

Exceptions

Iodine is lustrous.
Graphite conducts electricity despite being a non-metal.
Bromine is a liquid non-metal at room temperature.

4. Chemical Properties of Metals

Metals undergo several important chemical reactions.

4.1 Reaction of Metals with Oxygen

Most metals react with oxygen to produce metal oxides.

General reaction

Metal + Oxygen → Metal oxide

Example:

2Cu + O₂ → 2CuO

Copper forms black copper(II) oxide on heating.

Another example:

4Al + 3O₂ → 2Al₂O₃

Nature of metal oxides

Most metal oxides are basic.

However, some metal oxides show both acidic and basic behaviour. These are called amphoteric oxides.

Important examples:

Aluminium oxide – Al₂O₃
Zinc oxide – ZnO

Amphoteric oxides

An amphoteric oxide reacts with both acids and bases to form suitable products.

For aluminium oxide:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

and

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Reactivity with oxygen

Metals differ in their reactivity with oxygen.

Potassium and sodium react extremely vigorously.
Magnesium reacts less vigorously.
Copper does not burn easily but develops a black oxide coating when heated.
Gold and silver are very unreactive towards oxygen.

5. Reaction of Metals with Water

Different metals react with water differently.

General pattern

Metal + Water → Metal oxide + Hydrogen

If the metal oxide dissolves in water:

Metal oxide + Water → Metal hydroxide

Potassium and sodium

They react very violently with cold water and release hydrogen.

2K + 2H₂O → 2KOH + H₂ + heat

2Na + 2H₂O → 2NaOH + H₂ + heat

The reaction releases enough heat for the hydrogen produced to catch fire.

Calcium

Calcium reacts less violently with water:

Ca + 2H₂O → Ca(OH)₂ + H₂

Calcium may float because hydrogen bubbles stick to its surface.

Magnesium

Does not react significantly with cold water.
Reacts with hot water.
Produces magnesium hydroxide and hydrogen.

Aluminium, iron and zinc

Generally do not react with cold or hot water.
They can react with steam.

For example:

3Fe + 4H₂O → Fe₃O₄ + 4H₂

Lead, copper, silver and gold

These metals do not react with water under the conditions described in the chapter.

6. Reaction of Metals with Acids

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

General reaction

Metal + Dilute acid → Salt + Hydrogen

For example, a reactive metal such as iron reacts with dilute hydrochloric acid to form iron chloride and hydrogen.

Metals differ in their rates of reaction with acids.

The chapter gives the relative reactivity order:

Mg > Al > Zn > Fe

Copper does not react with dilute hydrochloric acid under the conditions discussed.

Nitric acid

Hydrogen is generally not released when metals react with nitric acid because nitric acid is a strong oxidising agent. It oxidises the hydrogen that would otherwise be produced.

An important exception mentioned in the chapter is that magnesium and manganese can release hydrogen with very dilute nitric acid.

7. Reaction with Solutions of Other Metal Salts

A more reactive metal can remove a less reactive metal from its salt solution.

General displacement reaction

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

This is called a displacement reaction.

Example:

Fe + CuSO₄ → FeSO₄ + Cu

Iron is more reactive than copper, so iron displaces copper from copper sulphate solution.

Key rule

A more reactive metal displaces a less reactive metal from its compound.

This principle is used to compare the reactivities of metals.

8. Reactivity Series

The reactivity series arranges metals from more reactive to less reactive.

Reactivity series

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

From top to bottom, reactivity decreases.

Important observations

Potassium is among the most reactive metals in the given series.
Gold is among the least reactive.
Metals above hydrogen can generally displace hydrogen from dilute acids.
A metal can displace another metal below it from an appropriate salt solution.

9. Why Metals and Non-metals React

Atoms tend to achieve a more stable electronic configuration.

Metals

Metals generally have relatively few electrons in their outermost shell.

They tend to lose electrons and form positively charged ions called cations.

Example:

Na → Na⁺ + e⁻

Non-metals

Non-metals generally tend to gain electrons to complete their outer shell.

Example:

Cl + e⁻ → Cl⁻

Thus:

Metals → lose electrons → positive ions.
Non-metals → gain electrons → negative ions.

10. Formation of Ionic Compounds

When electrons are transferred from a metal to a non-metal, oppositely charged ions are produced.

The electrostatic attraction between these ions forms an ionic compound.

Example: Sodium chloride

Sodium loses one electron:

Na → Na⁺ + e⁻

Chlorine gains one electron:

Cl + e⁻ → Cl⁻

The resulting Na⁺ and Cl⁻ ions attract one another and form NaCl.

Example: Magnesium chloride

Magnesium loses two electrons:

Mg → Mg²⁺ + 2e⁻

Each chlorine atom accepts one electron, so two chlorine atoms are needed.

Therefore:

MgCl₂

Compounds formed through electron transfer between metals and non-metals are called ionic compounds or electrovalent compounds.

11. Properties of Ionic Compounds

1. Physical state

Ionic compounds are generally:

Solid
Hard
Brittle

Their particles are held together by strong electrostatic attraction.

2. High melting and boiling points

They generally have high melting and boiling points because considerable energy is needed to overcome the attraction between ions.

3. Solubility

Many ionic compounds dissolve in water.

They are generally not soluble in organic liquids such as petrol and kerosene.

4. Electrical conductivity

Solid ionic compounds generally do not conduct electricity because their ions cannot move freely.
Molten ionic compounds conduct electricity.
Their aqueous solutions also conduct electricity because ions can move through the solution.

12. Occurrence of Metals

The Earth’s crust is a major natural source of metals.

Mineral

A naturally occurring element or compound found in the Earth’s crust is called a mineral.

Ore

A mineral from which a metal can be extracted profitably is called an ore.

Gangue

Ores often contain unwanted impurities such as soil and sand. These unwanted materials are called gangue.

13. Extraction of Metals

The method used to obtain a metal depends mainly on its position in the reactivity series.

Metals can broadly be divided into:

Low-reactivity metals
Medium-reactivity metals
Highly reactive metals

13.1 Low-reactivity Metals

Metals near the bottom of the activity series are relatively unreactive.

Their compounds may sometimes be converted into the metal simply by heating.

For example, mercury can be obtained from its ore through heating processes.

Copper can also be extracted from certain sulphide ores using suitable heating reactions.

13.2 Medium-reactivity Metals

Examples:

Zinc
Iron
Lead
Copper

These metals are often found as sulphides or carbonates.

Their ores are commonly first converted into oxides.

Roasting

Heating a sulphide ore strongly in the presence of excess air is called roasting.

Example:

2ZnS + 3O₂ → 2ZnO + 2SO₂

Calcination

Heating a carbonate ore strongly with limited air is called calcination.

Example:

ZnCO₃ → ZnO + CO₂

Reduction

The metal oxide is then reduced to obtain the metal.

For example:

ZnO + C → Zn + CO

Carbon acts as a reducing agent in this process.

14. Thermit Reaction

Some metal oxides can be reduced using a more reactive metal such as aluminium.

Example:

Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat

This reaction produces a very large amount of heat.

It is called the thermit reaction.

The heat produced can make iron molten, so this reaction has been used for joining railway tracks and repairing certain metal parts.

15. Extraction of Highly Reactive Metals

Highly reactive metals such as:

Potassium
Sodium
Calcium
Magnesium
Aluminium

cannot generally be extracted from their compounds by reducing them with carbon.

They have a stronger affinity for oxygen than carbon in the relevant extraction conditions.

Therefore, they are obtained through electrolytic reduction.

For example, sodium can be obtained by electrolysis of molten sodium chloride.

At the cathode:

Na⁺ + e⁻ → Na

At the anode:

2Cl⁻ → Cl₂ + 2e⁻

16. Refining of Metals

The metals obtained from ores are often not completely pure.

A common method of purification is electrolytic refining.

Electrolytic refining setup

Anode: Impure metal
Cathode: Thin strip of pure metal
Electrolyte: Solution containing a salt of the metal

When current is passed:

Metal from the impure anode enters the electrolyte.
Pure metal gets deposited on the cathode.
Soluble impurities remain in the solution.
Insoluble impurities settle below the anode as anode mud.

Copper, zinc, tin, nickel, silver and gold can be refined using electrolytic methods.

17. Corrosion

Corrosion is the gradual deterioration of a metal due to reactions with substances in its surroundings.

Rusting of iron

Iron rusts when it is exposed to both:

Air
Water/moisture

The brown coating formed on iron is called rust.

Other examples

Silver: develops a black coating due to formation of silver sulphide.

Copper: reacts with moist carbon dioxide and develops a green coating of basic copper carbonate.

18. Conditions Necessary for Rusting

The chapter’s experiment demonstrates that iron rusts when both air and water are available.

Air + water → rusting occurs.
Water without dissolved air → rusting does not occur under the experiment’s conditions.
Dry air without moisture → rusting does not occur.

Therefore, both oxygen from air and moisture are important for rusting of iron.

19. Prevention of Corrosion

Rusting can be prevented by protecting the metal from air and moisture.

Methods include:

Painting
Oiling
Greasing
Galvanising
Chrome plating
Anodising
Alloying

Galvanisation

Iron or steel is coated with a thin layer of zinc to protect it from rusting.

Anodising

A thicker protective oxide layer can be formed on aluminium by an electrochemical process. This improves its resistance to corrosion.

20. Alloys

An alloy is a homogeneous mixture of:

Two or more metals, or
A metal and a non-metal.

Alloying can improve useful properties such as strength, hardness and resistance to corrosion.

Examples

Brass

Copper + Zinc

Bronze

Copper + Tin

Solder

Lead + Tin
Has a relatively low melting point.

Stainless steel

Iron mixed with other elements including chromium and nickel.
It is hard and has good resistance to rusting.

Pure iron is relatively soft, so it is generally used in alloyed forms.

21. Gold and Alloys

Pure gold is very soft and is therefore not ideal for jewellery.

Gold is mixed with metals such as copper or silver to make it harder.

The chapter notes that 22-carat gold contains 22 parts pure gold mixed with 2 parts of another metal such as copper or silver.

22. Important Differences: Metals vs Non-metals

Property	Metals	Non-metals
Lustre	Generally lustrous	Generally dull
Malleability	Generally malleable	Generally non-malleable
Ductility	Generally ductile	Generally non-ductile
Conductivity	Usually good conductors	Usually poor conductors
Sonority	Generally sonorous	Generally non-sonorous
Ions formed	Usually positive ions	Usually negative ions
Oxides	Generally basic	Generally acidic or neutral
Reaction with dilute acids	Many metals release H₂	Generally do not displace H₂
Physical state	Mostly solids	Solids or gases; bromine is liquid

Remember: There are important exceptions, so these are general trends rather than absolute rules.

23. High-Value Chemical Equations

With oxygen

2Cu + O₂ → 2CuO

4Al + 3O₂ → 2Al₂O₃

With water

2Na + 2H₂O → 2NaOH + H₂ + heat

2K + 2H₂O → 2KOH + H₂ + heat

Ca + 2H₂O → Ca(OH)₂ + H₂

3Fe + 4H₂O → Fe₃O₄ + 4H₂

Amphoteric oxide

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Extraction

2ZnS + 3O₂ → 2ZnO + 2SO₂

ZnCO₃ → ZnO + CO₂

ZnO + C → Zn + CO

Thermit reaction

Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat

24. Most Important Definitions

Lustre: The shiny appearance of a metal surface.

Malleability: Ability of a substance to be beaten into thin sheets.

Ductility: Ability of a substance to be drawn into thin wires.

Sonorous: Property of producing a ringing sound when struck.

Amphoteric oxide: An oxide that reacts with both acids and bases.

Ionic compound: Compound formed through transfer of electrons and electrostatic attraction between oppositely charged ions.

Mineral: Naturally occurring element or compound found in the Earth’s crust.

Ore: A mineral from which a metal can be extracted profitably.

Gangue: Unwanted impurities associated with an ore.

Roasting: Strong heating of a sulphide ore in excess air.

Calcination: Strong heating of a carbonate ore in limited air.

Corrosion: Gradual deterioration of a metal due to environmental chemical reactions.

Alloy: Homogeneous mixture of two or more metals, or a metal and a non-metal.

Reactivity series: Arrangement of metals in decreasing order of their chemical reactivity.

Anode mud: Insoluble impurities that settle below the anode during electrolytic refining.

25. Exam-Focused Points to Memorise

Malleability → sheets
Ductility → wires
Sonorous → ringing sound
Graphite → non-metal that conducts electricity
Mercury → liquid metal at room temperature
Al₂O₃ and ZnO → amphoteric oxides
More reactive metal → displaces less reactive metal
Metals above H → generally displace H₂ from dilute acids
Roasting → sulphide ore + excess air
Calcination → carbonate ore + limited air
Highly reactive metals → electrolytic reduction
Impure metal → anode during electrolytic refining
Pure metal → cathode
Rusting → requires air and moisture
Galvanisation → zinc coating
Brass → Cu + Zn
Bronze → Cu + Sn
Solder → Pb + Sn
Thermit reaction → Fe₂O₃ + Al
Reactivity series:
K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

26. One-Page Quick Revision

Metals

Lustrous → Malleable → Ductile → Sonorous → Good conductors

Chemical behaviour

Metal + O₂ → Metal oxide

Metal + Water → Hydroxide/oxide + H₂

Metal + Dilute acid → Salt + H₂

More reactive metal + Salt → New salt + Less reactive metal

Reactivity

K Na Ca Mg Al Zn Fe Pb H Cu Hg Ag Au

Ionic compounds

Metal loses e⁻ → Cation

Non-metal gains e⁻ → Anion

Cation + Anion → Ionic compound

Metallurgy

Ore → Enrichment → Conversion to oxide → Reduction → Refining

Ore treatment

Sulphide → Roasting

Carbonate → Calcination

Corrosion

Iron + Air + Moisture → Rust

Protection

Paint / Oil / Grease / Galvanisation / Chrome plating / Anodising / Alloying

Alloys

Brass = Cu + Zn

Bronze = Cu + Sn

Solder = Pb + Sn

Stainless steel = Iron-based alloy with chromium and nickel

Chapter in One Sentence

Metals generally lose electrons, form positive ions and basic oxides, show characteristic reactions based on their position in the reactivity series, and are extracted and refined according to their reactivity; non-metals generally gain electrons and form negative ions and acidic or neutral oxides.