Notes :: Chapter 3: Metals and Non-metals

Physical Properties of Metals and Non-Metals

3.1.1 Metals

Activity 3.1: Metallic Lustre

Observation: Metals like iron, copper, aluminium, and magnesium have a characteristic shining surface. When cleaned with sandpaper, their lustre increases.
Conclusion: Metals, in their pure state, exhibit metallic lustre.

Activity 3.2: Hardness of Metals

Observation: Metals like iron, copper, aluminium, and magnesium are hard and cannot be easily cut with a knife. Sodium metal, however, is soft and can be easily cut with a knife.
Conclusion: Metals are generally hard, but their hardness varies. Sodium is an exception as it is soft.

Activity 3.3: Malleability

Observation: When hammered, some metals like iron, zinc, lead, and copper change shape and form thin sheets.
Conclusion: Malleability: The ability of metals to be beaten into thin sheets. Gold and silver are the most malleable metals.

Activity 3.4: Ductility

Observation: Metals like copper, aluminium, and iron can be drawn into wires. Gold is the most ductile metal – a 2 km long wire can be drawn from just 1 gram of gold.
Conclusion: Ductility: The ability of metals to be drawn into thin wires. This property allows metals to be used for electrical wires and other purposes.

Activity 3.5: Conductivity of Heat

Observation: A pin attached with wax to an aluminium or copper wire falls off when the wire is heated. This indicates that heat is transferred through the metal wire. Metals like silver and copper are the best conductors of heat. Lead and mercury are poor conductors of heat.
Conclusion: Metals are good conductors of heat. This property is why metals are used in cooking vessels.

Activity 3.6: Conductivity of Electricity

Observation: When a metal is placed in an electric circuit, the bulb glows, indicating the flow of electricity. Electric wires have a PVC or rubber coating for insulation.
Conclusion: Metals are good conductors of electricity. Insulating materials prevent electric shocks.

Sonority

Observation: When metals strike a hard surface, they produce a ringing sound.
Conclusion: Sonority: The ability of metals to produce a sound when struck. This property is why school bells are made of metals.

3.1.2 Non-Metals

Non-metals are fewer in number compared to metals. Examples: Carbon, Sulphur, Iodine, Oxygen, Hydrogen.

States of Matter: Non-metals are usually solids or gases. Exception: Bromine is the only liquid non-metal.

Activity 3.7: Physical Properties of Non-Metals

Observations from Activities 3.1 to 3.4 and 3.6 (compared with metals):

Element	Symbol	Type	Surface	Hardness	Malleability	Ductility	Conducts Electricity	Sonority
Carbon (Coal/Graphite)	C	Non-metal	Dull (Graphite is shiny)	Soft (except Diamond)	Brittle	Non-ductile	Graphite is a conductor	No
Sulphur	S	Non-metal	Dull	Soft	Brittle	Non-ductile	Non-conductor	No
Iodine	I	Non-metal	Lustrous	Soft	Brittle	Non-ductile	Non-conductor	No

Conclusions from Observations:

• Non-metals do not have metallic lustre (except iodine).
• They are generally soft and brittle.
• They are neither malleable nor ductile.
• They do not conduct electricity (except graphite).
• They are not sonorous (do not produce a ringing sound).

Exceptions to Physical Properties:

• Mercury is a metal but is liquid at room temperature.
• Gallium and Caesium have very low melting points and can melt in hand.
• Iodine, despite being a non-metal, is lustrous.

Carbon Allotropes:

• Diamond is the hardest natural substance and has a high melting point.
• Graphite is a good conductor of electricity, unlike most non-metals.
• Alkali metals (Lithium, Sodium, Potassium) are soft and can be cut with a knife.

Activity 3.8: Chemical Properties of Metals and Non-Metals

Experiment 1: Burning Magnesium (Metal)

Observation: Magnesium ribbon burns to form white ash (magnesium oxide). When dissolved in water, the solution turns red litmus paper blue.
Conclusion: Magnesium oxide is basic in nature.
2Mg + O₂ → 2MgO
MgO + H₂O → Mg(OH)₂ (Basic solution)

Experiment 2: Burning Sulphur (Non-Metal)

Observation: Sulphur burns to form sulphur dioxide (SO₂). When dissolved in water, the solution turns blue litmus paper red.
Conclusion: Sulphur dioxide forms an acidic solution (sulphurous acid).
S + O₂ → SO₂
SO₂ + H₂O → H₂SO₃ (Acidic solution)

Final Conclusion:

Metals form basic oxides, while non-metals form acidic oxides when dissolved in water. The classification of metals and non-metals is clearer based on their chemical properties rather than physical properties.

3.2 Chemical Properties of Metals

3.2.1 What Happens When Metals Are Burnt in Air?

Observation from Activity 3.9: Metals react with oxygen to form metal oxides.

General Reaction:
Metal + Oxygen → Metal Oxide

Example Reactions:

• Copper (Cu) burns to form black copper(II) oxide:
  2Cu + O₂ → 2CuO
• Aluminium (Al) forms aluminium oxide:
  4Al + 3O₂ → 2Al₂O₃

Types of Metal Oxides

• Basic Oxides: Most metal oxides react with acids to form salts and water.
• Amphoteric Oxides: Some metal oxides (e.g., aluminium oxide, zinc oxide) react with both acids and bases.
• Example:
  Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Example:
  Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (Sodium aluminate)

Solubility in Water

Most metal oxides are insoluble in water. Sodium oxide (Na₂O) and potassium oxide (K₂O) dissolve in water to form alkalis:

• Na₂O + H₂O → 2NaOH
• K₂O + H₂O → 2KOH

Reactivity of Metals with Oxygen

Highly reactive metals (e.g., sodium, potassium) react so vigorously that they catch fire when exposed to air. They are stored in kerosene to prevent accidental combustion.

Moderately reactive metals (e.g., magnesium, aluminium, zinc, lead) form a thin oxide layer that protects them from further oxidation.

Less reactive metals (e.g., iron) do not burn but iron filings burn vigorously.

Least reactive metals (e.g., copper, silver, gold) do not burn in air, but copper forms a black layer of CuO when heated.

Anodising Process

Anodising is a process to form a thick oxide layer on aluminium to improve corrosion resistance. It involves making an aluminium article the anode in an electrolytic cell containing dilute sulphuric acid. Oxygen gas evolved at the anode reacts with aluminium to make a thicker protective oxide layer, which can be dyed for aesthetic purposes.

Reactivity Order (Based on Oxygen Reaction)

Most reactive: Sodium (Na)
Moderate reactivity: Magnesium (Mg) > Aluminium (Al) > Zinc (Zn) > Iron (Fe)
Least reactive: Copper (Cu) > Silver (Ag) > Gold (Au)

3.2.2 What Happens When Metals React with Water?

Observation from Activity 3.10: Metals react with water to produce metal oxides and hydrogen gas. If the metal oxide is soluble in water, it forms a metal hydroxide.

General Reactions:

• Metal + Water → Metal Oxide + Hydrogen
• Metal Oxide + Water → Metal Hydroxide

Reactivity of Metals with Water

1. Highly Reactive Metals (React with Cold Water):
   Potassium (K) and Sodium (Na) react violently with cold water.
   The reaction is highly exothermic, causing hydrogen gas to catch fire.
   2K + 2H₂O → 2KOH + H₂ + heat
   2Na + 2H₂O → 2NaOH + H₂ + heat
2. Moderately Reactive Metals (React with Cold/Hot Water):
   Calcium (Ca) reacts less violently with cold water.
   The heat produced is not enough to ignite hydrogen.
   Ca + 2H₂O → Ca(OH)₂ + H₂
   Calcium starts floating due to hydrogen gas bubbles sticking to its surface.
   Magnesium (Mg) does not react with cold water but reacts with hot water, forming magnesium hydroxide and hydrogen gas.
   Mg + 2H₂O → Mg(OH)₂ + H₂
3. Less Reactive Metals (React with Steam Only):
   Aluminium (Al), Iron (Fe), and Zinc (Zn) do not react with cold or hot water but react with steam to form metal oxide and hydrogen gas.
   2Al + 3H₂O → Al₂O₃ + 3H₂
   3Fe + 4H₂O → Fe₃O₄ + 4H₂
4. Least Reactive Metals (No Reaction with Water):
   Lead (Pb), Copper (Cu), Silver (Ag), and Gold (Au) do not react with water at all.

Reactivity Order of Metals with Water

Most Reactive:

• Potassium (K) 🔥 (Violent reaction, catches fire)
• Sodium (Na) 🔥 (Violent reaction, catches fire)
• Calcium (Ca) (Reacts with cold water, hydrogen does not catch fire)
• Magnesium (Mg) (Reacts with hot water, not cold)
• Aluminium (Al), Zinc (Zn), Iron (Fe) (React with steam only)
• Lead (Pb), Copper (Cu), Silver (Ag), Gold (Au) (No reaction)

Least Reactive

The reactivity of metals decreases down the list. Highly reactive metals (K, Na) must be stored in kerosene to prevent reactions with moisture in the air.

3.2.3 What Happens When Metals React with Acids?

General Reaction:
Metals react with dilute acids to form salt and hydrogen gas.

Metal + Dilute Acid → Salt + Hydrogen gas

Observations from Activity 3.11:
Different metals react at different rates with dilute hydrochloric acid (HCl). More reactive metals react vigorously, producing bubbles of hydrogen gas rapidly. The temperature increase indicates the amount of heat released in the reaction.

Reactions of Some Metals with Dilute HCl

• Magnesium (Mg) reacts rapidly:
  Mg + 2HCl → MgCl₂ + H₂ ↑
  Fastest bubble formation, highest temperature rise.
• Aluminium (Al) reacts less vigorously than Mg:
  2Al + 6HCl → 2AlCl₃ + 3H₂ ↑
• Zinc (Zn) reacts moderately:
  Zn + 2HCl → ZnCl₂ + H₂ ↑
• Iron (Fe) reacts slowly:
  Fe + 2HCl → FeCl₂ + H₂ ↑
• Copper (Cu) does not react with dilute HCl:
  No bubbles were observed. No temperature change.

Reactivity Order with Dilute Acids

Mg > Al > Zn > Fe > Cu

Magnesium (Mg) is the most reactive. Copper (Cu) does not react at all.

Special Cases

Why Doesn’t Hydrogen Gas Evolve with Nitric Acid (HNO₃)?
Nitric acid (HNO₃) is a strong oxidizing agent. It oxidizes H₂ gas to water (H₂O), and itself gets reduced to nitrogen oxides (N₂O, NO, NO₂). Thus, no hydrogen gas is evolved in most cases.

Exceptions:
Magnesium (Mg) and Manganese (Mn) react with very dilute HNO₃ to produce H₂ gas.

Aqua Regia

Aqua regia is a freshly prepared mixture of:
3 parts concentrated hydrochloric acid (HCl)
1 part concentrated nitric acid (HNO₃)
It can dissolve gold (Au) and platinum (Pt), which neither acid can do alone. It is highly corrosive and fuming.

3.2.4 How Do Metals React with Solutions of Other Metal Salts?

Activity 3.12 Observations

• Copper wire in iron sulphate solution (FeSO₄):
  No visible change occurs. Copper does not react with FeSO₄ solution.
• Iron nail in copper sulphate solution (CuSO₄):
  The blue color of CuSO₄ solution fades. A reddish-brown deposit of copper forms on the iron nail. Reaction has occurred.

Explanation of Observations

Iron (Fe) is more reactive than copper (Cu). Iron displaces copper from copper sulphate (CuSO₄) solution, forming iron sulphate (FeSO₄) and depositing copper metal.

Reaction in the test tube with Fe and CuSO₄:
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Fe replaces Cu from CuSO₄ because Fe is more reactive than Cu.

Type of Reaction:

Displacement reaction.

No reaction in the test tube with Cu and FeSO₄ because:

Copper (Cu) is less reactive than iron (Fe). Cu cannot displace Fe from iron sulphate (FeSO₄) solution.

Conclusion from Activity 3.12

Iron (Fe) is more reactive than Copper (Cu). A more reactive metal can displace a less reactive metal from its salt solution. Reactivity series can be determined using displacement reactions.

Comparison with Previous Activities

Activity	Observation	Reactivity Order
3.9 (Reaction with Oxygen)	Sodium burns vigorously, Copper forms a black layer, Gold doesn’t react	Na > Mg > Zn > Fe > Cu > Ag
3.10 (Reaction with Water)	Sodium reacts explosively, Iron reacts with steam, Copper does not react	Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag
3.11 (Reaction with Acids)	Mg reacts most vigorously, Fe reacts slowly, Cu doesn’t react	Mg > Al > Zn > Fe > Cu
3.12 (Reaction with Metal Salts)	Fe displaces Cu from CuSO₄, Cu does not react with FeSO₄	Fe > Cu

Final Conclusion

The reactivity of metals can be determined by their ability to:

• Burn in oxygen (3.9)
• React with water (3.10)
• React with acids (3.11)
• Displace other metals from their salt solutions (3.12)

Iron (Fe) is more reactive than Copper (Cu) because it displaces Cu from its solution.

3.2.5 The Reactivity Series

Definition

The reactivity series is a list of metals arranged in decreasing order of their reactivity. The most reactive metals are at the top, while the least reactive ones are at the bottom.

Reactivity Series Table (Descending Order of Reactivity)

Metal	Reactivity
K (Potassium)	Most reactive
Na (Sodium)	Highly reactive
Ca (Calcium)	Very reactive
Mg (Magnesium)	Reactive
Al (Aluminium)	Moderately reactive
Zn (Zinc)	Less reactive
Fe (Iron)	Less reactive
Pb (Lead)	Poorly reactive
[H] (Hydrogen)	Reference (non-metal)
Cu (Copper)	Very less reactive
Hg (Mercury)	Very less reactive
Ag (Silver)	Least reactive
Au (Gold)	Least reactive

Significance of the Reactivity Series

Predicting Reactions

A more reactive metal can displace a less reactive metal from its compound. Example: Iron displaces copper from copper sulfate (Fe + CuSO₄ → FeSO₄ + Cu).

Reaction with Oxygen

Highly reactive metals (K, Na, Ca) react vigorously with oxygen. Less reactive metals (Fe, Pb, Cu) form a thin oxide layer. Unreactive metals (Ag, Au) do not react with oxygen.

Reaction with Water

Potassium & sodium react violently with cold water. Calcium reacts slowly with water. Iron & zinc react only with steam. Copper, silver, and gold do not react with water.

Reaction with Acids

Reactive metals (Mg, Zn, Fe) react with acids to release hydrogen gas. Less reactive metals (Cu, Ag, Au) do not react with acids.

Corrosion Resistance

Less reactive metals (Cu, Ag, Au) do not corrode easily and are used for jewelry & coins.

Conclusion

More reactive metals (like K, Na, Ca) are highly unstable and react easily . Moderately reactive metals (like Fe, Zn) react under specific conditions. Least reactive metals (like Ag, Au) are very stable and resistant to corrosion. Displacement reactions confirm the order of reactivity.

How Do Metals and Non-Metals React?

1. Reactivity of Elements and Electron Configuration

Noble gases have a completely filled valence shell, making them chemically inactive. Other elements react to achieve a stable electronic configuration (like noble gases).

2. Formation of Ions

Metals have 1, 2, or 3 electrons in their outermost shell and lose electrons to form cations (positively charged ions). Example: Sodium (Na) loses 1 electron to form Na⁺ (2,8). Magnesium (Mg) loses 2 electrons to form Mg²⁺ (2,8).

Non-Metals have 5, 6, or 7 electrons in their outermost shell and gain electrons to form anions (negatively charged ions). Example: Chlorine (Cl) gains 1 electron to form Cl⁻ (2,8,8).

3. Formation of Ionic Compounds (Electrovalent Compounds)

Metals lose electrons and Non-Metals gain electrons, forming oppositely charged ions. These ions are held together by strong electrostatic forces, forming an ionic bond.

4. Examples of Ionic Compounds

Sodium Chloride (NaCl)
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl

Magnesium Chloride (MgCl₂)
Mg → Mg²⁺ + 2e⁻
2Cl + 2e⁻ → 2Cl⁻
Mg²⁺ + 2Cl⁻ → MgCl₂

5. Cations and Anions in MgCl₂

Cation: Mg²⁺ (Magnesium ion)
Anion: Cl⁻ (Chloride ion)

Properties of Ionic Compounds

1. Physical Nature

Ionic compounds are solid and hard due to strong electrostatic forces between positive and negative ions. They are brittle and break into pieces under pressure.

2. Melting and Boiling Points

Ionic compounds have high melting and boiling points due to strong inter-ionic attractions that require a large amount of energy to break. Examples:
NaCl: Melting Point = 1074 K, Boiling Point = 1686 K
CaO: Melting Point = 2850 K, Boiling Point = 3120 K

3. Solubility

Soluble in water because water molecules weaken the electrostatic forces. Insoluble in organic solvents like petrol and kerosene.

4. Electrical Conductivity

Solid state: Ionic compounds do not conduct electricity because ions are fixed in position. Molten state or aqueous solution: They conduct electricity as ions become free to move and carry charge.

3.4 Occurrence and Extraction of Metals

3.4.1 Extraction of Metals

Metals in the Earth’s Crust:

• Least reactive metals (Low reactivity): Found in a free state (e.g., gold, silver, platinum, copper).
• Moderately reactive metals (Medium reactivity): Found in oxide, sulphide, or carbonate ores (e.g., zinc, iron, lead).
• Highly reactive metals (High reactivity): Never found free; always in combined form (e.g., potassium, sodium, calcium, magnesium, aluminum).

Reasons for the Abundance of Metal Oxides:

Oxygen is a highly reactive element and is abundant in nature. Many metals are found in the form of oxides due to their strong tendency to react with oxygen.

Three Categories of Metals Based on Reactivity:

• Low Reactivity Metals (e.g., Au, Ag, Cu): Found in native state (free state).
• Medium Reactivity Metals (e.g., Zn, Fe, Pb): Extracted using reduction with carbon.
• High Reactivity Metals (e.g., K, Na, Ca, Mg, Al): Extracted using electrolysis.

Steps in Metal Extraction from Ores:

1. Concentration of Ore: Removing impurities (gangue).
2. Reduction of Metal Compound: Extracting metal from its compound.
3. Refining: Purification of extracted metal.

3.4.2 Enrichment of Ores

Definition: Ores extracted from the earth contain impurities like soil, sand, and other unwanted materials, collectively called gangue. Enrichment of ores refers to the process of removing these impurities before metal extraction.

Methods of Enrichment:

The separation techniques used depend on the differences in physical or chemical properties between the ore and gangue. Some commonly used techniques include:

• Gravity Separation: Based on density differences.
• Magnetic Separation: Used when ore and gangue have different magnetic properties.
• Froth Flotation: Used for sulphide ores.
• Leaching: A chemical process to dissolve the ore selectively.

These processes help in obtaining a purified ore, which can then be used for extraction of metals.

Extracting Metals Low in the Activity Series

3.4.3 Extraction of Metals Low in the Activity Series

Characteristics: Metals low in the activity series (e.g., mercury, copper) are very unreactive. Their oxides can be reduced to metals by simple heating in air.

Example 1: Extraction of Mercury from Cinnabar (HgS)

1. Step 1: Roasting in air: Mercury sulphide (HgS) is heated in the presence of oxygen to form mercuric oxide (HgO).
   2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g)
2. Step 2: Further heating: Mercuric oxide (HgO) is heated to produce liquid mercury (Hg) and oxygen gas (O₂).
   2HgO(s) → 2Hg(l) + O₂(g)

Example 2: Extraction of Copper from Cu₂S

1. Step 1: Roasting in air: Copper sulphide (Cu₂S) reacts with oxygen to form copper oxide (Cu₂O) and sulphur dioxide (SO₂).
   2Cu₂S + 3O₂(g) → 2Cu₂O(s) + 2SO₂(g)
2. Step 2: Reduction of Copper Oxide: Copper oxide (Cu₂O) reacts with more copper sulphide (Cu₂S) to produce copper metal (Cu).
   2Cu₂O + Cu₂S → 6Cu(s) + SO₂(g)

Conclusion:

Metals low in the activity series are obtained by direct heating of their ores in air. No additional reducing agents are required for their extraction.

Extracting Metals in the Middle of the Activity Series

3.4.4 Extraction of Metals in the Middle of the Activity Series

Metals in this category: Iron (Fe), Zinc (Zn), Lead (Pb), Copper (Cu). These metals are moderately reactive and usually exist as sulphides or carbonates in nature.

Conversion of Ores into Oxides

Since it is easier to extract metals from their oxides than from sulphides or carbonates, the first step is conversion into oxides. This is done by two processes:

1. Roasting (for Sulphide Ores):
   Definition: Strong heating of sulphide ores in the presence of excess air to form metal oxides and sulphur dioxide (SO₂).
   Example:
   2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)
2. Calcination (for Carbonate Ores):
   Definition: Strong heating of carbonate ores in limited air to produce metal oxides and carbon dioxide (CO₂).
   Example:
   ZnCO₃(s) → ZnO(s) + CO₂(g)

Reduction of Metal Oxides

Once converted into oxides, metals are extracted by reducing agents such as carbon (coke) or highly reactive metals.

1. Reduction using Carbon (Coke):
   Metal oxides are heated with carbon, which reduces the oxide to metal and forms carbon monoxide (CO).
   Example:
   ZnO(s) + C(s) → Zn(s) + CO(g)
2. Reduction using More Reactive Metals (Displacement Reaction):
   Highly reactive metals like sodium (Na), calcium (Ca), aluminium (Al) are used to displace less reactive metals from their oxides.
   Example: Reduction of Manganese Oxide using Aluminium:
   3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + heat

Thermit Reaction (Iron Extraction and Welding)

Definition: A highly exothermic displacement reaction used for welding railway tracks or repairing machine parts.
Reaction:
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + heat
Iron (Fe) is obtained in the molten state due to the large amount of heat released.

Key Points

• Roasting is used for sulphide ores.
• Calcination is used for carbonate ores.
• Carbon (coke) or highly reactive metals are used as reducing agents.
• Thermit reaction is an example of displacement reaction used in metal extraction and welding.

Extraction of Highly Reactive Metals

Metals in this category:

Sodium (Na), Magnesium (Mg), Calcium (Ca), Aluminium (Al). These metals are very reactive and have high affinity for oxygen. Cannot be reduced using carbon because they form stronger bonds with oxygen than carbon does.

Method of Extraction: Electrolytic Reduction

Electrolysis of molten chlorides or oxides is used to extract these metals. Metals are deposited at the cathode (-ve electrode). Non-metals (like chlorine) are released at the anode (+ve electrode).

Example: Electrolysis of Sodium Chloride (NaCl) to Obtain Sodium (Na)

Reaction at Cathode (-ve electrode):
Na⁺ + e⁻ → Na

Reaction at Anode (+ve electrode):
2Cl⁻ → Cl₂ + 2e⁻

Sodium metal is deposited at the cathode, while chlorine gas is released at the anode.

Example: Electrolysis of Aluminium Oxide (Al₂O₃) to Obtain Aluminium (Al)

Aluminium is obtained by electrolytic reduction of aluminium oxide (bauxite purification process).

Key Points

• Highly reactive metals cannot be reduced using carbon.
• Electrolysis is used instead to extract these metals.
• Metals are deposited at the cathode (-ve electrode).
• Non-metals like chlorine (Cl₂) are released at the anode (+ve electrode).

3.4. 6 Refining of Metals

Metals obtained from reduction processes are not pure and contain impurities. Electrolytic Refining is the most common method to purify impure metals. Used for refining copper, zinc, tin, nickel, silver, gold, etc.

Electrolytic Refining Process

• Anode (-ve terminal): Impure metal is made the anode.
• Cathode (+ve terminal): A thin strip of pure metal is used as the cathode.
• Electrolyte: A solution of the metal’s salt is used as the electrolyte.

Process:
Current is passed through the electrolyte. Pure metal dissolves from the anode into the electrolyte. Equivalent pure metal is deposited on the cathode. Soluble impurities go into the solution. Insoluble impurities settle at the bottom as anode mud.

Key Points

• Electrolytic refining helps obtain pure metals.
• Impurities are removed as anode mud.
• The pure metal is deposited on the cathode.

Corrosion

Definition of Corrosion:

Corrosion is the process in which metals react with environmental factors such as air, water, or chemicals, leading to their deterioration.

Examples of Corrosion:

• Silver tarnishing: Silver reacts with sulphur in the air, forming a black coating of silver sulphide (Ag₂S).
• Copper corrosion: Copper reacts with moist carbon dioxide in the air, forming a green layer of basic copper carbonate (CuCO₃·Cu(OH)₂).
• Rusting of iron: Iron reacts with moist air, forming brown flaky rust (Fe₂O₃·xH₂O).

Activity 3.14 - Rusting of Iron:

Objective: To study the conditions required for rusting.

Procedure: Take three test tubes labeled A, B, and C and place clean iron nails in each.

• Test Tube A: Contains water and air → Rusting occurs.
• Test Tube B: Contains boiled distilled water + oil layer (prevents air contact) → No rusting.
• Test Tube C: Contains anhydrous calcium chloride (absorbs moisture, keeps air dry) → No rusting.

Conclusion:

Rusting occurs only when both air (oxygen) and water (moisture) are present.

Prevention of rusting:

• Keeping iron dry
• Using protective coatings (e.g., painting, galvanization)
• Using rust-resistant alloys (e.g., stainless steel)

Methods to Prevent Rusting of Iron:

• Painting: Prevents contact with air and moisture.
• Oiling and Greasing: Creates a protective layer against moisture.
• Galvanization: Coating iron/steel with zinc to prevent rusting. Even if the zinc layer is broken, it protects the underlying metal by acting as a sacrificial layer.
• Chrome Plating: A layer of chromium prevents oxidation and corrosion.
• Anodizing: Enhances corrosion resistance, commonly used for aluminium.
• Alloying: Mixing metals to improve properties and prevent rusting.

Alloying and Its Benefits:

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

Process: The primary metal is melted, mixed with other elements in definite proportions, and then cooled.

Examples of Alloys and Their Uses:

Alloy	Composition	Properties & Uses
Stainless Steel	Iron + Nickel + Chromium	Hard, rust-resistant, used in utensils, tools
Brass	Copper + Zinc	Low conductivity, used in decorations & musical instruments
Bronze	Copper + Tin	Corrosion-resistant, used in statues & medals
Solder	Lead + Tin	Low melting point, used in welding electrical wires
Gold Alloy (22 Carat Gold)	22 parts Gold + 2 parts Silver/Copper	Harder than pure gold, used in jewelry

Special Case – The Iron Pillar of Delhi:

Built over 1600 years ago by Indian iron workers. Rust-resistant due to advanced ancient metallurgy. 8 meters high and weighs 6000 kg (6 tonnes). Attracts global scientific study for its corrosion resistance.