Chapter 3: Metals and Non-Metals - Vardaan Master Notes

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

Elements are classified broadly into Metals and Non-metals based on their physical and chemical properties. Let's examine their fundamental characteristics through key laboratory investigations:

Key Observations: Fundamental Physical Properties

Metallic Lustre: Metals in their pure state possess a shining surface.
Hardness: Metals are generally hard, though hardness varies from metal to metal.
Exception: Alkali metals (Li, Na, K) are so soft they can be cut with a knife.
Malleability: Metals can be beaten into thin sheets without breaking. Gold and Silver are the most malleable metals.
Ductility: The ability of metals to be drawn into thin wires. Gold is the most ductile of all metals (1g can be drawn into a 2km wire!).

Experimental Investigation: Thermal Conductivity

Experimental Setup: 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).

Experimental Investigation: 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.

Comparative Analysis: Non-Metals & Their Oxides

Non-Metals: Examples include Carbon, Sulphur, Iodine, and Oxygen. They exist mostly as solids or gases at room temperature (with the sole exception of Bromine, which is liquid).

Nature of Oxides:
• Most non-metals produce acidic oxides when dissolved in water (e.g., $SO_2$).
• Most metals give rise to 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 (I₂) 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)

Key Practical 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

Experimental Investigation: Reaction with Water & Steam

Experimental Setup: Action of steam on a metal sample

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.

Experimental Investigation: Metal Displacement Reactions

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 (Formation of Ionic Compounds)

Elements react with each other to achieve a stable, completely filled valence shell (Noble gas octet configuration).

Metals have 1, 2, or 3 electrons in their valence shell. They lose these electrons readily to form positive ions (cations).
Non-metals have 4 to 7 electrons in their valence shell. They gain electrons to complete their octet and form negative ions (anions).
The compounds formed by the transfer of electrons from a metal to a non-metal are called Ionic Compounds or Electrovalent Compounds.

Core Visual Mechanism

Electron Transfer Mechanism — Formation of NaCl & MgCl₂

Formation of Sodium Chloride (NaCl)

$$Na\; (2,8,1) \longrightarrow Na^+\; (2,8) + e^-\quad (\text{Sodium cation})$$ $$Cl\; (2,8,7) + e^- \longrightarrow Cl^-\; (2,8,8)\quad (\text{Chloride anion})$$

Crucial Fact: Sodium chloride does NOT exist as discrete single molecules, but as large 3D aggregates of oppositely charged ions held firmly by intense electrostatic attraction forces.

Formation of Magnesium Chloride (MgCl₂)

$$Mg\; (2,8,2) \longrightarrow Mg^{2+}\; (2,8) + 2e^-\quad (\text{Magnesium cation})$$ $$2Cl\; (2,8,7) + 2e^- \longrightarrow 2Cl^-\; (2,8,8)\quad (\text{Two Chloride anions})$$

Magnesium loses 2 valence electrons; one electron is transferred to each of the two chlorine atoms to achieve complete noble gas octets.

General 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. Occurrence of Metals (Extraction of Metals)

The earth's crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc.

Essential Metallurgical Definitions

Minerals: The elements or chemical compounds which occur naturally in the earth's crust are known as minerals.
Ores: At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

Key Conceptual Principle

Why are the ores of many metals oxides?

Oxygen is a very reactive element and is very abundant on the earth. Therefore, many metals react with atmospheric oxygen over time to form stable oxide ores.

Principles of Metallurgy (Extraction of Metals)

Some metals are found in the earth's crust in the free state, while some are found in the form of their compounds:

Low in Activity Series: The metals at the bottom of the activity series are the least reactive and are often found in a free state (e.g., gold, silver, platinum, and copper). Copper and silver are also found in the combined state as their sulphide or oxide ores.
Top of Activity Series: The metals at the top of the activity series (K, Na, Ca, Mg, and Al) are so reactive that they are never found in nature as free elements.
Middle of Activity Series: The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive and are found in the earth's crust mainly as oxides, sulphides, or carbonates.

Flowchart: Major steps involved in the extraction of metals from ores

Summary: Extraction Methods Based on Activity Series

Reactivity Category	Metals	Nature of Ore	Method of Extraction
Metals of High Reactivity	K, Na, Ca, Mg, Al	Molten Ore (Chlorides/Oxides)	Electrolysis of molten ore (cannot be reduced by carbon).
Metals of Medium Reactivity	Zn, Fe, Pb, Cu	• Carbonate ore • Sulphide ore	• Calcination (for carbonates) • Roasting (for sulphides) → Oxide of metal → Reduction to metal using Carbon → Purification.
Metals of Low Reactivity	Ag, Au (free state) Hg, Cu (sulphides)	Sulphide ores or Native state	Roasting → Metal → Refining (reduction by heat alone).

Step 1: Enrichment of Ores (Concentration)

Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue.

The impurities must be removed from the ore prior to the extraction of the metal. The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. Different separation techniques are accordingly employed.

Step 2A: Extracting Metals Low in the Activity Series

Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone.

Representative Extraction Reactions (Low Reactivity)

1. Extraction of Mercury from Cinnabar (HgS):

Cinnabar ($HgS$) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide ($HgO$). Mercuric oxide is then reduced to mercury on further heating:

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

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

2. Extraction of Copper from Copper Sulphide (Cu₂S):

Similarly, copper which is found as $Cu_2S$ in nature can be obtained from its ore by just heating in air:

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

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

Step 2B: Extracting Metals in the Middle of the Activity Series

The metals in the middle of the activity series such as iron, zinc, lead, copper, are moderately reactive. These are usually present as sulphides or carbonates in nature.

Key Chemical Rationale

Why convert sulphides and carbonates into oxides before reduction?

It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides.

Roasting vs. Calcination

Process	Roasting	Calcination
Type of Ore	Used for Sulphide ores (e.g., Zinc blende ZnS).	Used for Carbonate ores (e.g., Calamine ZnCO₃).
Heating Condition	Heating strongly in the presence of excess air.	Heating strongly in limited air.
Reaction Equation	$$2ZnS(s) + 3O_2(g) \overset{\text{Heat}}{\longrightarrow} 2ZnO(s) + 2SO_2(g)$$	$$ZnCO_3(s) \overset{\text{Heat}}{\longrightarrow} ZnO(s) + CO_2(g)$$

Reduction of Metal Oxides to Metals:

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc:

$$ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$$

Note: Obtaining metals from their compounds is also a reduction process.

Displacement Reactions as Reduction:

Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc., are used as reducing agents because they can displace metals of lower reactivity from their compounds.

For example, when manganese dioxide is heated with aluminium powder, the following reaction takes place:

$$3MnO_2(s) + 4Al(s) \rightarrow 3Mn(l) + 2Al_2O_3(s) + \text{Heat}$$

The Thermit Process (Exothermic Reduction)

These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals are produced in the molten state.

In fact, the reaction of iron(III) oxide ($Fe_2O_3$) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermit reaction:

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

Substances Oxidised & Reduced: Aluminium ($Al$) is oxidised (gains oxygen to form $Al_2O_3$), while Iron(III) oxide ($Fe_2O_3$) is reduced (loses oxygen to form $Fe$).

Step 2C: Extracting Metals High in the Activity Series

The metals high up in the reactivity series (K, Na, Ca, Mg, Al) are very reactive. They cannot be obtained from their compounds by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals have more affinity for oxygen than carbon.

Electrolytic Reduction of Molten Chlorides

These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides.

The metals are deposited at the cathode (the negatively charged electrode):
$$\mathbf{Na^+ + e^- \rightarrow Na}$$
Chlorine is liberated at the anode (the positively charged electrode):
$$\mathbf{2Cl^- \rightarrow Cl_2 + 2e^-}$$
Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.

Step 3: Refining of Impure Metals

The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining (used for copper, zinc, tin, nickel, silver, gold, etc.).

Industrial Process: Electrolytic Refining of Copper

Experimental Setup: Electrolytic refining of copper in acidified copper sulphate bath

Anode: Impure copper.
Cathode: A strip of pure copper.
Electrolyte: A solution of acidified copper sulphate.
Process: On passing electric current through the electrolyte, pure copper from the anode dissolves into the electrolyte. An equivalent amount of pure copper from the electrolyte is deposited on the cathode.
Fate of Impurities: The soluble impurities go into the solution, whereas the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

Concept Mastery Check & Essential Questions

Define the following terms: (i) Mineral, (ii) Ore, (iii) Gangue.
Answer:
(i) Mineral: The elements or compounds which occur naturally in the earth's crust are known as minerals.
(ii) Ore: Minerals that contain a very high percentage of a particular metal and from which the metal can be profitably extracted are called ores.
(iii) Gangue: Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue.

Name two metals which are found in nature in the free state.
Answer: Gold and Platinum (Silver and Copper are also found in the free state).

What chemical process is used for obtaining a metal from its oxide?
Answer: Reduction process. Metal oxides are reduced to metals by heating with reducing agents like carbon (for moderately reactive metals like zinc), displacement using reactive metals like aluminium (thermit process), or electrolytic reduction (for highly reactive metals like sodium and aluminium).

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).

Real-World 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.

Laboratory Investigation: Conditions Necessary for Rusting

Experimental Setup: Testing the rusting of iron nails under varying atmospheric conditions

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 ≈ 80%) + Zinc (Zn ≈ 20%)	Malleable, lustrous golden finish, lower electrical conductivity than pure Cu. Used for musical instruments, decorative items, and fittings.
Bronze	Copper (Cu ≈ 90%) + Tin (Sn ≈ 10%)	Tough, resistant to corrosion, poor conductor of electricity. Used for statues, coins, and medals.
Solder	Lead (Pb ≈ 50%) + Tin (Sn ≈ 50%)	Low melting point (lower than constituent metals). Used for welding electrical wires together.
Amalgam	Mercury (Hg) + Any other metal (e.g., Na or Ag)	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.

Metals and Non-Metals