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 ($Hg$) is liquid at room temperature.
Melting Point	Gallium ($Ga$) and Caesium ($Cs$) melt on your palm (very low melting points).
Lustre	Iodine ($I_2$) is a non-metal but it is lustrous with a shining surface.
Allotropy (Carbon)	Diamond (hardest natural substance, high MP) & Graphite (good electrical 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

The earth's crust is the major source of metals. Soluble salts of metals are also present in seawater (such as sodium chloride and magnesium chloride). Metals occur either in the free state (low reactive metals like Gold, Silver, Platinum) or in the form of compounds (oxides, sulphides, carbonates).

Fundamental Definitions

Mineral: Naturally occurring elements or chemical compounds found in the earth's crust.
Ore: Minerals containing a high percentage of a particular metal from which the metal can be extracted profitably and conveniently. (All ores are minerals, but not all minerals are ores).
Gangue (Matrix): Unwanted rocky, sandy, or earthy impurities associated with naturally mined ores.
Metallurgy: The entire scientific and technological process used for extracting pure metals from their ores.

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

1. Enrichment of Ores (Concentration)

Ores mined from the earth are associated with large amounts of gangue (sand, soil, clay). These impurities must be removed prior to the extraction of the metal. The concentration process is based on the physical and chemical differences between the gangue and the ore (e.g., hydraulic washing, magnetic separation, froth floatation, or chemical leaching).

2. Extracting Metals Low in the Activity Series (Low Reactivity)

Metals at the bottom of the activity series are unreactive. Oxides of these metals can be reduced to metals by heat alone without requiring any additional chemical reducing agent.

Low Reactivity Extractions

Example 1: Extraction of Mercury from Cinnabar ($HgS$):

When cinnabar (mercuric sulphide) is heated strongly in air, it first converts to mercuric oxide ($HgO$), which on further heating decomposes into liquid mercury:

$2HgS(s) + 3O_2(g) \xrightarrow{\text{Heat}} 2HgO(s) + 2SO_2(g)$

$2HgO(s) \xrightarrow{\text{Heat}} 2Hg(l) + O_2(g)$

Example 2: Extraction of Copper from Copper Glance ($Cu_2S$):

Copper sulphide ore is heated in air to convert a part of it into copper(I) oxide. The remaining $Cu_2S$ then reacts with $Cu_2O$ (auto-reduction / self-reduction) to yield copper metal:

$2Cu_2S(s) + 3O_2(g) \xrightarrow{\text{Heat}} 2Cu_2O(s) + 2SO_2(g)$

$2Cu_2O(s) + Cu_2S(s) \xrightarrow{\text{Heat}} 6Cu(s) + SO_2(g)$

3. Extracting Metals in the Middle of the Activity Series (Medium Reactivity)

Metals in the middle of the activity series (such as $Fe, Zn, Pb, Cu$) usually occur as sulphide or carbonate ores in nature. It is much easier to obtain a metal from its oxide than directly from its sulphide or carbonate. Therefore, ores are first converted into metal oxides by Roasting or Calcination.

Comparison: Roasting vs. Calcination

Feature	Roasting	Calcination
Type of Ore	Used for Sulphide ores (e.g., $ZnS, PbS, Cu_2S$).	Used for Carbonate ores & hydrated oxides (e.g., $ZnCO_3, CaCO_3$).
Air Supply	Heated strongly in excess air (oxygen) below melting point.	Heated strongly in limited air or in the absence of air.
Gas Released	Evolves pungent Sulphur dioxide ($SO_2$) gas.	Evolves colourless Carbon dioxide ($CO_2$) gas.
Chemical Equation	$$2ZnS + 3O_2 \xrightarrow{\text{Heat}} 2ZnO + 2SO_2 \uparrow$$	$$ZnCO_3 \xrightarrow{\text{Heat}} ZnO + CO_2 \uparrow$$

4. Reduction of Metal Oxides to Metals

Once converted to oxides, metals are obtained through reduction:

A. Reduction using Carbon (Coke / Smelting):

Metal oxides are heated with a suitable reducing agent such as carbon (coke):

$ZnO(s) + C(s) \xrightarrow{\text{Heat}} Zn(s) + CO(g)$

$Fe_2O_3(s) + 3C(s) \xrightarrow{\text{Heat}} 2Fe(s) + 3CO(g)$

B. Reduction using Displacement Reactions (Aluminothermy / Thermit Process):

Highly reactive metals like Aluminium ($Al$), Sodium ($Na$), or Calcium ($Ca$) can be used as reducing agents because they displace metals of lower reactivity from their oxides. These displacement reactions are highly exothermic, producing the extracted metal in a molten liquid state.

Thermit Reaction (Welding Railway Tracks)

The reaction of iron(III) oxide ($Fe_2O_3$) with aluminium powder is known as the Thermit Reaction:

$$Fe_2O_3(s) + 2Al(s) \xrightarrow{\text{Ignition}} 2Fe(l) + Al_2O_3(s) + \text{Huge amount of Heat}$$

Application: The molten iron produced is directly used for joining cracked railway tracks and heavy broken machine parts.

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

5. Extracting Metals Towards the Top of the Activity Series (High Reactivity)

Metals high in the reactivity series ($K, Na, Ca, Mg, Al$) cannot be reduced by carbon because these metals have a much greater affinity for oxygen than carbon does. Therefore, they are extracted by Electrolytic Reduction of their molten chlorides or oxides.

Electrolytic Reduction of Molten NaCl

When electricity is passed through molten sodium chloride ($NaCl$):

At Cathode (Negative Electrode): Sodium ions gain electrons (reduction) to deposit as sodium metal:
$$\mathbf{Na^+ + e^- \rightarrow Na(s)}$$
At Anode (Positive Electrode): Chloride ions lose electrons (oxidation) to liberate chlorine gas:
$$\mathbf{2Cl^- \rightarrow Cl_2(g) + 2e^-}$$
Note: Aluminium is similarly extracted by the electrolytic reduction of molten aluminium oxide (alumina, $Al_2O_3$).

6. Refining of Metals: Electrolytic Refining

Metals produced by reduction processes contain impurities and must be refined. The most widely used method for refining metals like Copper, Zinc, Tin, Nickel, Silver, and Gold is Electrolytic Refining.

Activity 3.14: Electrolytic Refining of Copper

Fig 4.2: Electrolytic refining of copper apparatus showing cathode, anode, and anode mud

Anode (+): Thick slab of impure copper metal.
Cathode (-): Thin strip of pure copper metal.
Electrolyte: Acidified aqueous solution of copper sulphate ($CuSO_4$).
Reaction Mechanism:
- At Anode: $Cu\text{ (impure)} \rightarrow Cu^{2+} + 2e^-$ (Copper dissolves into electrolyte).
- At Cathode: $Cu^{2+} + 2e^- \rightarrow Cu\text{ (pure)}$ (Pure copper deposits onto cathode strip).
Impurities Fate: Soluble impurities go directly into solution, while insoluble impurities (like gold, silver, platinum) settle at the bottom of the anode as Anode Mud.

5. Corrosion and its Prevention

Corrosion is the slow eating away or deterioration of metals by chemical or electrochemical reaction with atmospheric substances (moisture, oxygen, carbon dioxide, hydrogen sulphide).

NCERT Case Studies of Corrosion

Silver: Articles turn black after some time when exposed to air because silver reacts with traces of hydrogen sulphide ($H_2S$) / sulphur in the air to form a black coating of Silver Sulphide ($Ag_2S$).
Copper: Reacts with moist carbon dioxide ($CO_2$) and water in the air to slowly lose its shiny brown surface and form a green protective coat of Basic Copper Carbonate [$CuCO_3 \cdot Cu(OH)_2$].
Iron: Acquires a reddish-brown flaky coating called Rust (Hydrated Iron(III) oxide, $Fe_2O_3 \cdot xH_2O$) when exposed to moist air for a prolonged time.

Activity 3.14: Investigating Conditions Necessary for Rusting

Fig 5.1: Investigating the conditions under which iron rusts in Test Tubes A, B, and C

Three Test Tubes Setup:

Test Tube A: Contains iron nails dipped in water and exposed to air. Result: Nails rust heavily because both air and water are present.
Test Tube B: Contains boiled distilled water covered with a layer of oil to prevent air from dissolving. Result: Nails do not rust (air is absent).
Test Tube C: Contains dry air and anhydrous calcium chloride ($CaCl_2$) to absorb all moisture. Result: Nails do not rust (moisture is absent).
Conclusion: Both oxygen (air) and moisture (water) are essential for the rusting of iron.

Prevention of Corrosion

Rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising, or by making alloys.

Galvanisation

Galvanisation: A method of protecting iron and steel from rusting by coating them with a thin layer of Zinc ($Zn$).

Why is it effective? Zinc is more reactive than iron. Even if the zinc coating is scratched or broken, zinc oxidises preferentially (sacrificial protection), preventing the underlying iron from rusting.

Alloying: Improving Metal Properties

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by melting the primary metal and then dissolving the other elements in definite proportions, followed by cooling to room temperature.

Important Alloys Reference Table

Alloy	Composition	Key Characteristics & Primary Uses
Steel	Iron ($Fe$) + Carbon ($0.05\% - 1.5\%$)	Pure iron is very soft and stretches easily when hot; adding small amount of carbon makes it hard and strong.
Stainless Steel	Iron ($Fe$) + Nickel ($Ni$) + Chromium ($Cr$)	Hard, highly ductile, and does not rust. Used for surgical instruments and utensils.
Brass	Copper ($Cu \approx 80\%$) + Zinc ($Zn \approx 20\%$)	Malleable, lustrous golden finish, lower electrical conductivity than pure Cu. Used for musical instruments, decorative items, and fittings.
Bronze	Copper ($Cu \approx 90\%$) + Tin ($Sn \approx 10\%$)	Tough, resistant to corrosion, poor conductor of electricity. Used for statues, coins, and medals.
Solder	Lead ($Pb \approx 50\%$) + Tin ($Sn \approx 50\%$)	Low melting point (lower than constituent metals). Used for welding electrical wires together.
Amalgam	Mercury ($Hg$) + Any metal (e.g., Sodium or Silver)	Dental amalgam ($Ag-Sn-Hg$) used for tooth fillings; Sodium amalgam ($Na-Hg$) used as a reducing agent.

The Iron Pillar at Delhi (Ancient Indian Metallurgy)

The famous Iron Pillar near the Qutub Minar in Delhi was built more than 1600 years ago by Indian iron workers. It is 8 metres high and weighs approximately 6 tonnes (6000 kg). Despite centuries of exposure to sun and rain, it has not rusted due to the formation of a protective thin passive layer of magnetic iron oxide ($\mathrm{Fe_3O_4}$) on its surface, standing as testament to ancient India's advanced metallurgical prowess.

The Gold Standard (24-Carat vs 22-Carat Gold)

Pure gold, known as 24 Carat gold, is very soft and pliable, making it unsuitable for making sturdy jewelry. To improve hardness, it is alloyed with either Silver ($Ag$) or Copper ($Cu$). In India, 22 Carat gold is commonly used, which means 22 parts by mass of pure gold is alloyed with 2 parts of either copper or silver.