METALS AND NON-METALS

Chapter 03 | High-Fidelity Board Study Module

1. Physical Properties (Metals vs. Non-Metals)

In Class IX, you learned that elements are classified into Metals and Non-metals based on their properties. Let's examine these through NCERT-mandated activities.

Activity 3.1 - 3.4: The Basics

Lustre (3.1): Metals in their pure state have a shining surface. This is called Metallic Lustre.
Hardness (3.2): Metals are generally hard, but the hardness varies.
Exception: Alkali metals (Li, Na, K) are so soft they can be cut with a knife.
Malleability (3.3): Metals can be beaten into thin sheets. Gold and Silver are the most malleable metals.
Ductility (3.4): The ability of metals to be drawn into thin wires. Gold is the most ductile (1g can make a 2km wire!).

Activity 3.5: Thermal Conductivity

Fig 3.5: Action of heat on a metal wire (Testing Thermal Conductivity)

Experiment: Heat one end of a metal wire with a pin attached via wax at the other end.

Observation: The wax melts and the pin falls off, but the metal wire does not melt.

Inference: Metals are good conductors of heat and have high melting points.
Best Conductors: Silver (Ag) and Copper (Cu).
Poor Conductors: Lead (Pb) and Mercury (Hg).

Activity 3.6: Electrical Conductivity

Experiment: Set up a circuit with a gap for testing metal samples.

Observation: The bulb glows when metals like Cu or Al are connected.

Safety Note: Electric wires in homes are coated with PVC (Polyvinylchloride) or a rubber-like material because they are insulators.

Activity 3.7 - 3.8: Non-Metals

Non-Metals (3.7): Examples include Carbon, Sulphur, Iodine, Oxygen. They are either solids or gases (except Bromine which is liquid).

Nature of Oxides (3.8):
• Most Non-metals produce Acidic Oxides (e.g., $SO_2$).
• Deep-Cut: Some non-metal oxides are Neutral (e.g., Carbon Monoxide ($CO$), Nitric Oxide ($NO$), Nitrous Oxide ($N_2O$)).
• Most Metals produce Basic Oxides (e.g., $MgO$).

Board Favorite: Exceptions Table

Property	Exception Detail
State	Mercury is liquid at room temperature.
Melting Point	Gallium and Caesium melt on your palm (very low MP).
Lustre	Iodine is a non-metal but it is lustrous.
Allotropy (Carbon)	Diamond (Hardest natural substance) & Graphite (Good conductor).

Exam Definition Sonorous: Metals produce a ringing sound when struck hard. This is why school bells are made of metals.

2. Chemical Properties of Metals

Metals behave differently when they react with air, water, and other substances. Their reactivity is the basis of the Reactivity Series.

A. What happens when Metals are burnt in Air?

Almost all metals combine with oxygen to form Metal Oxides. Most are basic, but some show unique properties.

Metal + Oxygen $\rightarrow$ Metal Oxide

Example: $2Cu + O_2 \rightarrow 2CuO$ (Black Copper(II) Oxide)

Example: $4Al + 3O_2 \rightarrow 2Al_2O_3$ (Aluminium Oxide)

NCERT Observation: Iron does not burn on heating, but iron filings burn vigorously when sprinkled in the flame of the burner. Similarly, Copper does not burn, but is coated with a thin layer of black copper(II) oxide.

Board Favorite: Amphoteric Oxides

Oxides which react with both acids and bases to produce salt and water are called Amphoteric Oxides.

Example 1 (Aluminium):

$Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$
$Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 \text{ (Sodium Aluminate)} + H_2O$

Example 2 (Zinc):

$ZnO + 2HCl \rightarrow ZnCl_2 + H_2O$
$ZnO + 2NaOH \rightarrow Na_2ZnO_2 \text{ (Sodium Zincate)} + H_2O$

Process: Anodising

Anodising is a process of forming a thick oxide layer of aluminium. This layer makes it resistant to further corrosion. During the process, clean Al is made the Anode and electrolysed with dilute sulphuric acid. Oxygen gas evolved reacts with Al to form a protective oxide layer.

Note: This layer can be dyed easily to give aluminium articles an attractive finish.

B. What happens when Metals react with Water?

Metal + Water $\rightarrow$ Metal Oxide + Hydrogen gas
Metal Oxide + Water $\rightarrow$ Metal Hydroxide

Activity 3.10: The Water Test

Fig 3.10: Action of steam on a metal

Na & K: React violently with Cold Water. The evolved $H_2$ immediately catches fire (Exothermic).
Calcium (Ca): Reacts less violently. It floats because bubbles of $H_2$ stick to its surface.
Magnesium (Mg): Does not react with cold water; reacts with Hot Water to form $Mg(OH)_2$ and floats.
Al, Fe, Zn: React only with Steam to form Metal Oxides (not hydroxides).
$2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$
$3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$
Pb, Cu, Ag, Au: Do not react with water at all.

C. Reaction with Acids (The $HNO_3$ Exception)

Generally, Metal + Dilute Acid $\rightarrow$ Salt + $H_2$. However, Nitric Acid ($HNO_3$) is different.

THE $HNO_3$ LOGIC

Hydrogen gas is NOT evolved when a metal reacts with $HNO_3$ because it is a strong oxidising agent. It oxidises the $H_2$ produced to $H_2O$ and itself gets reduced to nitrogen oxides ($N_2O, NO, NO_2$).

Exceptions: Magnesium ($Mg$) and Manganese ($Mn$) react with very dilute $HNO_3$ to evolve $H_2$ gas.

Fact: Aqua Regia (Royal Water)

A freshly prepared mixture of Concentrated HCl and Concentrated $HNO_3$ in the ratio 3:1. It is a highly corrosive, fuming liquid that can dissolve Gold and Platinum, even though neither acid can do so alone.

3. Reactivity Series & Ionic Bonding

Why do some metals react vigorously while others remain inert? The answer lies in their position in the Reactivity Series.

Activity 3.12: Metal Displacement

Metal A + Salt Solution of B $\rightarrow$ Salt Solution of A + Metal B

If Metal A displaces Metal B from its solution, it is more reactive than B. This is the most reliable way to compare reactivities.

The Reactivity Series

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

(Most Reactive $\rightarrow$ Potassium | Least Reactive $\rightarrow$ Gold)

Reaction of Metals with Non-Metals

Elements react to achieve a stable, completely filled valence shell (Noble gas configuration).
Electronic Logic: Sodium (2,8,1) loses 1e- to become $Na^+$ (2,8 - Neon config). Chlorine (2,8,7) gains 1e- to become $Cl^-$ (2,8,8 - Argon config).

Ionic Bond Formation

Formation of $MgCl_2$:
$Mg (2,8,2) \rightarrow Mg^{2+} (2,8) + 2e^-$
$2Cl (2,8,7) + 2e^- \rightarrow 2Cl^- (2,8,8)$
$Mg^{2+} + [:\ddot{Cl}:]^- \times 2 \rightarrow MgCl_2$

Ionic Compounds: Formed by the transfer of electrons from a metal to a non-metal. These are also called Electrovalent Compounds.

Activity 3.13: Properties of Ionic Compounds

Physical Nature: Solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. They are generally brittle.
Melting & Boiling Points: They have high MP and BP. Reason: A considerable amount of energy is required to break the strong inter-ionic attraction.
Solubility: Generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
Conduction of Electricity:
• Solid State: Do NOT conduct (ions cannot move).
• Molten/Aqueous State: Conduct electricity as ions move freely.

4. Metallurgy: Extraction of Metals

Earth's crust is the major source of metals. Some occur in the free state (Gold, Silver, Platinum), while most are found as compounds.

Definitions

Mineral: Elements or compounds occurring naturally in the earth's crust.
Ore: Minerals from which a metal can be extracted profitably.
Gangue: Large amounts of impurities like soil and sand found in ores.

Fig 4.1: Summary flowchart for extraction of metals from their ores

A. Metals at the Bottom (Low Reactivity)

Ores are often Sulphides. They are reduced to metals by heating in air (Roasting).

Example 1: Cinnabar ($HgS$):
$2HgS + 3O_2 \xrightarrow{\text{Heat}} 2HgO + 2SO_2$
$2HgO \xrightarrow{\text{Heat}} 2Hg + O_2$

Example 2: Copper Glance ($Cu_2S$):
$2Cu_2S + 3O_2 \xrightarrow{\text{Heat}} 2Cu_2O + 2SO_2$
$2Cu_2O + Cu_2S \xrightarrow{\text{Heat}} 6Cu + SO_2$

B. Metals in the Middle (Medium Reactivity)

Usually present as Sulphides or Carbonates. It is easier to extract metals from their oxides, so they are first converted:

Roasting: Heating sulphide ores strongly in excess air.
$2ZnS + 3O_2 \xrightarrow{\text{Heat}} 2ZnO + 2SO_2$
Calcination: Heating carbonate ores strongly in limited air.
$ZnCO_3 \xrightarrow{\text{Heat}} ZnO + CO_2$

Reduction: Metal oxides are then reduced using Carbon (Coke) or displacement with more reactive metals like Aluminium.

THERMIT REACTION

The reaction of $Fe_2O_3$ with Aluminium is highly exothermic; the iron is produced in a molten state.

$Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$

Sister Reaction: Manganese dioxide with Aluminium:
$3MnO_2(s) + 4Al(s) \rightarrow 3Mn(l) + 2Al_2O_3(s) + \text{Heat}$

Application: Used to join railway tracks or cracked machine parts.

C. Metals at the Top (High Reactivity)

They cannot be reduced by carbon because they have more affinity for oxygen than carbon. They are extracted by Electrolytic Reduction.

Example: Sodium is obtained by electrolysis of molten sodium chloride.
• At Cathode (-): $Na^+ + e^- \rightarrow Na$
• At Anode (+): $2Cl^- \rightarrow Cl_2 + 2e^-$

Activity 3.14: Electrolytic Refining

Fig 4.2: Electrolytic refining of copper

Process for Copper:

Anode (+): Impure metal.
Cathode (-): Thin strip of pure metal.
Electrolyte: Solution of the metal salt (Acidified $CuSO_4$).
Result: Pure metal from anode dissolves into electrolyte and deposits onto cathode. Insoluble impurities settle as Anode Mud.

5. Corrosion and its Prevention

Corrosion is the gradual destruction of metals by chemical reaction with their environment (Air + Moisture).

NCERT Case Studies

Silver: Becomes black when exposed to air because it reacts with Sulphur to form a coating of Silver Sulphide ($Ag_2S$).
Copper: Reacts with moist Carbon Dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is Basic Copper Carbonate.
Iron: When exposed to moist air for a long time, it acquires a coating of a brown flaky substance called Rust.

Activity 3.14: Conditions for Rusting

Fig 5.1: Investigating the conditions under which iron rusts

Experiment: Test iron nails in three conditions.

Result: Rusting only occurs in Tube A where both Air and Water are present.
Prevention: Painting, oiling, greasing, galvanising, chrome plating, or making alloys.

Definition: Galvanisation

A method of protecting steel and iron from rusting by coating them with a thin layer of Zinc. The galvanised article is protected even if the zinc coating is broken.

Alloying: Improving Metal Properties

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.

High-Value Alloy List

Alloy	Composition	Key Property
Steel	Fe + C (0.05%)	Hard and strong.
Stainless Steel	Fe + Ni + Cr	Hard and does not rust.
Brass	Cu + Zn	Electrical conductivity lower than pure Cu.
Bronze	Cu + Sn	Poor conductor of electricity.
Solder	Pb + Sn	Low melting point (for welding).
Amalgam	Metal + Mercury	Variable.

More to Know: The Iron Pillar

The Iron Pillar near the Qutub Minar in Delhi was built more than 1600 years ago. It is 8 metres high and weighs approximately 6 tonnes (6000 kg). It has resisted rusting for centuries due to the formation of a thin film of magnetic oxide ($Fe_3O_4$) on its surface, showcasing ancient India's metallurgy.

The Gold Standard

Pure gold (24 Carat) is very soft and not suitable for making jewelry. It is alloyed with either Silver or Copper to make it hard. Generally, in India, 22 Carat gold is used (22 parts gold + 2 parts Cu/Ag).