Vardaan Learning Institute

Two most malleable and ductile metals: Gold ($\text{Au}$) and Silver ($\text{Ag}$) are the most malleable and ductile metals known. (A single gram of gold can be drawn into a wire 2 kilometres long).

1. Physical State: Mercury ($\text{Hg}$) is a liquid at room temperature, whereas all other metals are solids.
2. Hardness: Alkali metals (Sodium $\text{Na}$, Potassium $\text{K}$, Lithium $\text{Li}$) are extremely soft solids that can be easily sliced with a kitchen knife.
3. Melting & Boiling Points: Gallium ($\text{Ga}$) and Caesium ($\text{Cs}$) have exceptionally low melting points (29.7°C and 28.4°C). They melt if placed on a human palm.
4. Density: Sodium and potassium have low densities (less than water, $< 1\text{ g/cm}^3$), meaning they float on water.

Reason for coating: Polyvinyl chloride (PVC) or rubber-like materials are covalent organic polymers. They are electrical insulators (do not conduct electricity). Coating the metallic wires with insulating PVC ensures safety, preventing short circuits and protecting users from electric shocks.

Property	Metals	Non-metals
(a) Hardness	Generally hard (except alkali metals).	Generally soft (except diamond).
(b) State	Solids at room temp (except liquid mercury).	Can be solids, liquids (bromine), or gases.
(c) Lustre	Have metallic lustre; can be polished.	Dull and do not reflect light (except iodine).
(d) Conductivity	Excellent thermal and electrical conductors.	Poor conductors (insulators, except graphite).
(e) Density	High density and high melting/boiling points.	Low density and low melting/boiling points.

1. Iodine ($\text{I}_2$): It is a non-metal, yet it has a shiny, metallic, lustrous appearance.
2. Carbon in the form of Diamond: It is an allotrope of carbon (non-metal), yet it is the hardest natural substance known and has an extremely high melting point (~3550°C).
3. Carbon in the form of Graphite: It is a non-metal allotrope, yet it is a good conductor of electricity due to the presence of free delocalised electrons in its layered structure.
4. Bromine ($\text{Br}_2$): It is the only non-metal that exists as a liquid at room temperature.

1. Silver vs Copper wiring: Silver ($\text{Ag}$) has the lowest electrical resistivity and is the absolute best conductor. However, copper ($\text{Cu}$) is used because silver is extremely expensive, rare, and lacks high tensile strength. Copper is highly ductile, strong, abundant, and cheap, making it the most viable material for household wiring.
2. Cutting Sodium/Potassium: Alkali metals have weak metallic bonding due to the presence of only a single valence electron per atom and large atomic sizes. This results in a very soft crystal lattice that easily yields under physical cutting pressure.

1. Sonorous metals: The property of emitting a deep, ringing sound when struck with a hard object.
   Practical use: Metals like bronze and brass are used to make temple/church bells and musical instruments (cymbals, gongs) because of this ringing resonance.
2. Metallic lustre: The clean, bright, reflective surface of metals in their pure state.
   Practical use: Highly reflective gold, silver, and platinum are used to make fine jewellery and decorative plating. Aluminium and silver coatings are used on mirrors to reflect light.

Reason: In diamond, each carbon atom is covalently bonded to four other carbon atoms in a rigid, three-dimensional tetrahedral giant covalent network. Breaking this massive network of strong covalent bonds requires an enormous amount of thermal energy, leading to its exceptionally high melting and boiling points.

Property demonstrated: This demonstrates the physical property of malleability (ability to be hammered into sheets without fracturing).

When a solid piece of sulphur is struck with a hammer, it shatters instantly into a fine powder or small, irregular fragments.
Explanation: Solid non-metals are held together by weak intermolecular van der Waals forces. When a shearing stress is applied, the layers of atoms shift and easily break apart since there are no metallic bonds or sea of mobile electrons to hold the lattice together under distortion.

Balanced chemical equation: $$2\text{Cu}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{CuO}(s)$$ Change in appearance: The reddish-brown copper metal is coated with a dull, black layer of copper(II) oxide ($\text{CuO}$).

Why immersed in kerosene? Sodium does not react with hydrocarbons like kerosene. Immersing it in kerosene cuts off all contact with atmospheric oxygen and moisture, preventing accidental combustion.
If dropped in water: It reacts violently, swimming on water, releasing hydrogen gas which immediately catches fire due to the high heat liberated: $$2\text{Na}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq) + \text{H}_2(g)\uparrow + \text{Heat}$$ $$4\text{Na}(s) + \text{O}_2(g) \rightarrow 2\text{Na}_2\text{O}(s) \text{ (Reaction with oxygen)}$$

Examples: Aluminium oxide ($\text{Al}_2\text{O}_3$) and Zinc oxide ($\text{ZnO}$).
Balanced equations for Aluminium oxide ($\text{Al}_2\text{O}_3$):

1. Reaction with acid ($\text{HCl}$): $$\text{Al}_2\text{O}_3(s) + 6\text{HCl}(aq) \rightarrow 2\text{AlCl}_3(aq) + 3\text{H}_2\text{O}(l)$$
2. Reaction with base ($\text{NaOH}$): $$\text{Al}_2\text{O}_3(s) + 2\text{NaOH}(aq) \rightarrow 2\text{NaAlO}_2(aq) + \text{H}_2\text{O}(l)$$ (Product is Sodium aluminate).

Balanced equations for Zinc oxide ($\text{ZnO}$):

1. Reaction with acid: $$\text{ZnO}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2\text{O}(l)$$
2. Reaction with base: $$\text{ZnO}(s) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{ZnO}_2(aq) + \text{H}_2\text{O}(l)$_ (Product is Sodium zincate).

1. Sodium reacts violently: Sodium combines immediately and violently with cold water, releasing hydrogen gas and forming sodium hydroxide. The reaction is so exothermic that the hydrogen instantly catches fire and burns with a yellow flame: $$2\text{Na}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaOH}(aq) + \text{H}_2(g)\uparrow$$
2. Calcium floats: Calcium reacts less violently with cold water to form calcium hydroxide and hydrogen gas: $$\text{Ca}(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{H}_2(g)\uparrow$$ Reason for floating: The bubbles of hydrogen gas produced during the reaction stick to the surface of the calcium metal, carrying the metal piece to the top, where it floats.
3. Iron reacts only with steam: Iron does not react with cold or hot water. It only reacts when steam is passed over heated iron filings, producing iron(II,III) oxide and hydrogen gas: $$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \text{ (Steam)} \rightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)\uparrow$$

Equation for Magnesium with HCl: $$\text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g)\uparrow$$ Test for gas: Collect the gas and bring a burning splint near the mouth of the tube. The gas will burn with a sharp 'pop' sound, indicating it is hydrogen.

When a metal reacts with $\text{HNO}_3$, hydrogen gas ($\text{H}_2$) is initially produced. However, the strong oxidizing $\text{HNO}_3$ immediately oxidizes this hydrogen gas to water ($\text{H}_2\text{O}$). In the process, the nitric acid itself is reduced to any of the nitrogen oxides (such as nitric oxide $\text{NO}$, nitrogen dioxide $\text{NO}_2$, or nitrous oxide $\text{N}_2\text{O}$).
Exception: Very dilute $\text{HNO}_3$ reacts with magnesium ($\text{Mg}$) and manganese ($\text{Mn}$) to evolve hydrogen gas: $$\text{Mg}(s) + 2\text{HNO}_3(dilute) \rightarrow \text{Mg(NO}_3)_2(aq) + \text{H}_2(g)\uparrow$$

Composition: A freshly prepared mixture of concentrated Hydrochloric acid ($\text{HCl}$) and concentrated Nitric acid ($\text{HNO}_3$) in a 3:1 volume ratio.
Why it dissolves Gold/Platinum: Although neither single concentrated acid reacts with gold, the combination reacts to generate nascent chlorine ($\text{Cl}$): $$3\text{HCl}(conc.) + \text{HNO}_3(conc.) \rightarrow \text{NOCl}(aq) + 2\text{H}_2\text{O}(l) + 2\text{Cl}(g)$$ This highly reactive nascent chlorine attacks the gold atoms, forming soluble auric chloride ($\text{AuCl}_3$), which dissolves completely: $$\text{Au}(s) + 3\text{Cl}(g) \rightarrow \text{AuCl}_3(aq)$$

Iron nail in Copper Sulphate: When an iron nail ($\text{Fe}$) is immersed in blue copper sulphate ($\text{CuSO}_4$) solution, iron, being more reactive than copper, displaces copper from its salt solution (a single displacement reaction): $$\text{Fe}(s) + \text{CuSO}_4(aq) \text{ (Blue)} \rightarrow \text{FeSO}_4(aq) \text{ (Pale Green)} + \text{Cu}(s) \text{ (Red-brown)}$$ Observations:

1. The blue colour of the copper sulphate solution gradually fades and turns pale green due to the formation of iron(II) sulphate ($\text{FeSO}_4$).
2. A reddish-brown deposit of metallic copper is formed on the iron nail.

1. $$2\text{Al}(s) + 3\text{H}_2\text{O}(g) \rightarrow \text{Al}_2\text{O}_3(s) + 3\text{H}_2(g)\uparrow$$
2. $$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \rightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)\uparrow$$
3. $$\text{Ca}(s) + 2\text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{H}_2(g)\uparrow$$

• Copper with dilute HCl: In the reactivity series, copper lies below hydrogen. This means copper is less reactive than hydrogen and cannot reduce hydronium ions from dilute $\text{HCl}$ to evolve hydrogen gas. Hence, no reaction occurs.
• Copper with hot concentrated H2SO4: Hot concentrated sulphuric acid is a strong oxidizing agent (unlike dilute acid). It reacts with copper not by hydrogen displacement, but by oxidizing copper to copper sulphate, being itself reduced to sulphur dioxide gas ($\text{SO}_2$): $$\text{Cu}(s) + 2\text{H}_2\text{SO}_4(conc.) \xrightarrow{\Delta} \text{CuSO}_4(aq) + 2\text{H}_2\text{O}(l) + \text{SO}_2(g)\uparrow$$

Balanced chemical equation: $$\text{Cu}(s) + 2\text{AgNO}_3(aq) \rightarrow \text{Cu(NO}_3)_2(aq) + 2\text{Ag}(s)\downarrow$$ Observations:

1. The colourless silver nitrate solution gradually turns a beautiful blue due to the formation of hydrated copper(II) ions [$\text{Cu(NO}_3)_2$].
2. Shiny, needle-like crystals of silver metal deposit on the copper strip.

Decrease in reactivity:

• Sodium ($\text{Na}$): Reacts violently at room temperature, catching fire instantly. Must be stored in kerosene. $$4\text{Na}(s) + \text{O}_2(g) \rightarrow 2\text{Na}_2\text{O}(s)$$
• Magnesium ($\text{Mg}$): Does not react at room temperature. It must be heated, burning with a dazzling white flame. $$2\text{Mg}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{MgO}(s)$$
• Zinc ($\text{Zn}$): Burns only on strong heating. $$2\text{Zn}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{ZnO}(s)$$
• Iron ($\text{Fe}$): Does not burn even on heating, but iron filings burn vigorously when sprinkled in the flame. $$3\text{Fe}(s) + 2\text{O}_2(g) \xrightarrow{\Delta} \text{Fe}_3\text{O}_4(s)$$
• Copper ($\text{Cu}$): Does not burn at all. It only forms a black surface layer of oxide on prolonged strong heating. $$2\text{Cu}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{CuO}(s)$$

Why noble gases are inert: Noble gases (Helium $\text{He}$, Neon $\text{Ne}$, Argon $\text{Ar}$, etc.) already possess a completely filled outermost shell (a stable duplet for He, and stable octet of 8 electrons for others). Because their valence shell is completely filled, they have zero chemical drive to gain, lose, or share electrons, making them highly inert.

Ionic bond: The chemical bond formed by the complete transfer of one or more electrons from a metal atom to a non-metal atom, resulting in strong electrostatic forces of attraction holding the oppositely charged ions together.

Sodium (atomic number 11, electronic configuration 2, 8, 1) has 1 valence electron. It loses this electron to form a sodium cation ($\text{Na}^+$): $$\text{Na} \rightarrow \text{Na}^+ + e^- \text{ (2, 8, 1 } \rightarrow \text{ 2, 8)}$$ Chlorine (atomic number 17, configuration 2, 8, 7) has 7 valence electrons. It gains the electron lost by sodium to form a chloride anion ($\text{Cl}^-$): $$\text{Cl} + e^- \rightarrow \text{Cl}^- \text{ (2, 8, 7 } \rightarrow \text{ 2, 8, 8)}$$ Electron-dot transfer representation: $$\text{Na}^{\bullet} + \cdot \ddot{\text{C}}\text{l}: \rightarrow [\text{Na}]^+ [:\ddot{\text{C}}\text{l}:]^- \rightarrow \text{NaCl}$$

Magnesium (atomic number 12, configuration 2, 8, 2) has 2 valence electrons. It must lose both electrons to achieve a stable octet, forming a magnesium cation ($\text{Mg}^{2+}$): $$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \text{ (2, 8, 2 } \rightarrow \text{ 2, 8)}$$ Each chlorine atom (2, 8, 7) needs only 1 electron to complete its octet: $$\text{Cl} + e^- \rightarrow \text{Cl}^-$$ Why two chlorine atoms are required: Since a magnesium atom loses 2 electrons but one chlorine atom can only accept 1 electron, two chlorine atoms are required to accept the 2 electrons lost by a single magnesium atom.
Electron-dot transfer representation: $$\text{Mg}^{\bullet\bullet} + 2 \cdot\ddot{\text{C}}\text{l}: \rightarrow [\text{Mg}]^{2+} \left[ 2 \,\, : \ddot{\text{C}}\text{l}: \right]^- \rightarrow \text{MgCl}_2$$

1. Formation of Magnesium Oxide ($\text{MgO}$): $$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^- \text{ (2, 8, 2 } \rightarrow \text{ 2, 8)}$$ $$\text{O} + 2e^- \rightarrow \text{O}^{2-} \text{ (2, 6 } \rightarrow \text{ 2, 8)}$$ $$\text{Mg}^{\bullet\bullet} + \cdot\ddot{\text{O}}\cdot \rightarrow [\text{Mg}]^{2+} [:\ddot{\text{O}}:]^{2-} \rightarrow \text{MgO}$$
2. Formation of Calcium Oxide ($\text{CaO}$): $$\text{Ca} \rightarrow \text{Ca}^{2+} + 2e^- \text{ (2, 8, 8, 2 } \rightarrow \text{ 2, 8, 8)}$$ $$\text{O} + 2e^- \rightarrow \text{O}^{2-} \text{ (2, 6 } \rightarrow \text{ 2, 8)}$$ $$\text{Ca}^{\bullet\bullet} + \cdot\ddot{\text{O}}\cdot \rightarrow [\text{Ca}]^{2+} [:\ddot{\text{O}}:]^{2-} \rightarrow \text{CaO}$$

1. Physical Nature: They are hard, crystalline, brittle solids due to strong electrostatic forces holding the ions in a rigid crystal lattice.
2. Melting and Boiling Points: They have exceptionally high melting and boiling points.
3. Solubility: They are generally soluble in polar solvents like water, but completely insoluble in non-polar organic solvents (such as kerosene, benzene, ether).
4. Electrical Conductivity: They do not conduct electricity in the solid state, but are highly conductive in the molten state or when dissolved in water.

Reason: Overcoming these highly stable electrostatic attractions to melt or boil the compound requires a massive amount of thermal energy to break the rigid crystal lattice. Therefore, ionic compounds exhibit exceptionally high melting and boiling points (e.g., $\text{NaCl}$ melts at 801°C).

• In solid state: Ions are tightly locked in fixed positions in the rigid 3D crystal lattice by strong electrostatic forces. Since the ions are not free to move, solid ionic compounds cannot conduct electricity.
• In molten/aqueous state: Dissolving in water (due to polar water molecules breaking the lattice) or heating to a melt provides enough energy to overcome the electrostatic attractions. The crystal lattice breaks down, and the ions become free-moving mobile charge carriers. Thus, they conduct electricity efficiently.

1. Mineral: Naturally occurring inorganic homogeneous substances containing a metal in a free or combined state found in the Earth's crust.
2. Ore: Those specific minerals from which a metal can be extracted profitably, easily, and commercially on a large scale.
3. Gangue: The unwanted rocky, earthy, sandy, and siliceous impurities associated with ores when mined from the Earth.

Difference: "All ores are minerals, but all minerals are not ores." For example, clay contains aluminium, but it is not an ore because extraction is unprofitable. Bauxite is both a mineral and an ore because aluminium can be extracted from it profitably.

Three major steps:

1. Concentration of Ore (Enrichment): Removal of gangue impurities.
2. Conversion to Metal Oxide and Reduction: Converting the ore to oxide and reducing it to crude metal.
3. Refining: Purifying the crude metal to obtain pure metal.

Two physical processes:

1. Hydraulic Washing (Gravity Separation): Separation based on difference in densities of ore and gangue.
2. Magnetic Separation: Separation based on magnetic properties of ore or gangue.
3. Froth Flotation (used for sulphide ores): Separation based on differential wetting properties of ore and gangue with oil and water.

Extraction of Mercury from Cinnabar ($\text{HgS}$): First, cinnabar is heated in air (roasting) to convert it into mercuric oxide ($\text{HgO}$): $$2\text{HgS}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{HgO}(s) + 2\text{SO}_2(g)$$ On further heating, mercuric oxide decomposes directly into liquid mercury metal and oxygen gas: $$2\text{HgO}(s) \xrightarrow{\text{Strong Heating}} 2\text{Hg}(l) + \text{O}_2(g)$$

First, the ore is heated in oxygen, partially converting it into copper(I) oxide ($\text{Cu}_2\text{O}$): $$2\text{Cu}_2\text{S}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{Cu}_2\text{O}(s) + 2\text{SO}_2(g)$$ When the supply of air is reduced, the formed copper(I) oxide reacts with the remaining copper(I) sulphide to undergo self-reduction, producing blister copper: $$2\text{Cu}_2\text{O}(s) + \text{Cu}_2\text{S}(s) \xrightarrow{\Delta} 6\text{Cu}(l) + \text{SO}_2(g)\uparrow$$

Property	Roasting	Calcination
Ore Type	Used primarily for sulphide ores.	Used primarily for carbonate or hydrated ores.
Air Supply	Heated strongly in the presence of excess air.	Heated strongly in the absence or limited air.
Gas Evolved	Sulphur dioxide gas ($\text{SO}_2$) is evolved.	Carbon dioxide gas ($\text{CO}_2$) or water vapour is evolved.
Equation Example	$$2\text{ZnS}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{ZnO}(s) + 2\text{SO}_2(g)$$	$$\text{ZnCO}_3(s) \xrightarrow{\Delta} \text{ZnO}(s) + \text{CO}_2(g)$$

Balanced chemical equation (ZnO with carbon): $$\text{ZnO}(s) + \text{C}(s) \xrightarrow{\Delta} \text{Zn}(g)\uparrow + \text{CO}(g)$$ In this reaction, zinc oxide is reduced to zinc metal, and carbon is oxidized to carbon monoxide.

Balanced chemical equation: $$\text{Fe}_2\text{O}_3(s) + 2\text{Al}(s) \rightarrow 2\text{Fe}(l) + \text{Al}_2\text{O}_3(s) + \text{Heat}$$ Industrial application: Because a massive amount of heat is evolved, the iron metal is produced in a molten/liquid state. It is used for welding railway tracks and filling cracks in heavy machinery parts on-site.

Extraction method: Electrolytic reduction of their molten salts.
Electrolysis of molten Sodium Chloride: When molten $\text{NaCl}$ is electrolysed, sodium ions are reduced at the cathode, and chloride ions are oxidised at the anode:

• At the Cathode (-): Sodium ions gain electrons and deposit as liquid sodium metal. $$\text{Na}^+ + e^- \rightarrow \text{Na}(l)$$
• At the Anode (+): Chloride ions lose electrons to form chlorine gas. $$2\text{Cl}^- \rightarrow \text{Cl}_2(g)\uparrow + 2e^-$$

It is the purification of crude copper using electrolysis.

1. Anode: A thick block of impure crude copper.
2. Cathode: A thin strip of pure copper.
3. Electrolyte: Acidified copper sulphate ($\text{CuSO}_4$) solution.

Process: When current is passed, copper ions dissolve from the anode into the electrolyte, while an equivalent amount of pure copper ions from the electrolyte deposit onto the cathode.
Electrode reactions:

• At the Anode (+) (Oxidation): $$\text{Cu}(s) \text{ (Impure)} \rightarrow \text{Cu}^{2+}(aq) + 2e^-$$
• At the Cathode (-) (Reduction): $$\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s) \text{ (Pure)}$$

Why it forms: Impurities that are less reactive than copper (lower in the reactivity series) do not dissolve in the acidic copper sulphate bath because they cannot be oxidized easily at the operating potential. Instead, they fall off as the anode dissolves.
Substances found in it: Valuable noble metals such as Gold ($\text{Au}$), Silver ($\text{Ag}$), Platinum ($\text{Pt}$), and Selenium. Recovering these noble metals offsets the electrical costs of the refining process.

$$\text{Concentration of ore} \rightarrow \text{Roasting/Calcination} \rightarrow \text{Reduction} \rightarrow \text{Refining}$$ Explanation of order:

1. Concentration of ore: First, we must remove gangue to reduce bulk and prevent side-reactions.
2. Roasting/Calcination: Metals must be converted into their oxides because oxides are much easier to reduce than carbonates or sulphides.
3. Reduction: The metal oxide is then reduced to obtain crude metal.
4. Refining: Finally, the crude metal contains impurities and must be purified to obtain highly pure metal suitable for commercial applications.

Three metals and changes:

1. Iron ($\text{Fe}$): Rusts in moist air, forming a flaky, reddish-brown coating.
   Chemical formula: $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$ (Hydrated iron(III) oxide).
2. Silver ($\text{Ag}$): Tarnishes, forming a dull, black coating when exposed to air containing trace hydrogen sulphide gas.
   Chemical formula: $\text{Ag}_2\text{S}$ (Silver sulphide).
3. Copper ($\text{Cu}$): Corrodes in moist air containing carbon dioxide, forming a green coating.
   Chemical formula: $\text{CuCO}_3 \cdot \text{Cu(OH)}_2$ (Basic copper carbonate).

Experimental proof (Three test tubes):

1. Test Tube A (Both air and water): Clean iron nails are placed in normal tap water with air in the headspace.
   Observation: Nails rust heavily, proving both are needed.
2. Test Tube B (Only water, no air): Nails are placed in boiled distilled water (boiling drives out dissolved air), and a layer of oil is poured on top to prevent air from dissolving back.
   Observation: Nails do not rust, proving air is necessary.
3. Test Tube C (Only air, no water): Nails are placed in a tube containing anhydrous calcium chloride ($\text{CaCl}_2$), which absorbs all atmospheric moisture.
   Observation: Nails do not rust, proving moisture is necessary.

How it protects even if scratched: Zinc is more reactive than iron. When the zinc coating is scratched, exposing the iron underneath, zinc acts as a sacrificial anode. Zinc has a higher tendency to lose electrons than iron: $$\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$$ Zinc oxidizes and corrodes in preference to the iron. The electrons released by zinc flow to the iron, preventing the iron from losing electrons and rusting. This is called cathodic/sacrificial protection.

1. Painting, greasing, or oiling: Creates a physical barrier on the metal surface, blocking all contact with atmospheric oxygen and moisture. Oiling/greasing is preferred for moving machinery parts since it also reduces friction.
2. Tinning and chrome plating: Coating iron with a layer of unreactive, non-corroding metals like tin ($\text{Sn}$) or chromium ($\text{Cr}$) via electroplating. It provides a shiny protective layer. Tin is safe for food cans.
3. Alloying: Mixing the metal homogeneously with other metals or non-metals. This alters the crystal lattice structure, significantly reducing chemical reactivity and making the resulting alloy highly resistant to corrosion (e.g., mixing iron with carbon, chromium, and nickel to form rust-proof stainless steel).

Preparation: Prepared by melting the primary metal, dissolving the other elements in it in definite proportions, and cooling the mixture to room temperature.
Constituents, properties, and uses:

1. Brass:
   • Constituents: Copper ($\text{Cu}$, ~70%) and Zinc ($\text{Zn}$, ~30%).
   • Properties: Malleable, beautiful golden appearance, corrosion-resistant, conducts electricity.
   • Uses: Making cooking utensils, musical instruments, decorative items, and ammunition shells.
2. Bronze:
   • Constituents: Copper ($\text{Cu}$, ~90%) and Tin ($\text{Sn}$, ~10%).
   • Properties: Extremely tough, highly resistant to corrosion, sonorous.
   • Uses: Making statues, medals, coins, bells, and machine bearings.
3. Solder:
   • Constituents: Lead ($\text{Pb}$, ~50%) and Tin ($\text{Sn}$, ~50%).
   • Properties: Has an exceptionally low melting point (lower than both constituent metals), solidifies quickly.
   • Uses: Electrical soldering to join wires and copper connections in electronics.

Stainless Steel:

• Constituents: Iron ($\text{Fe}$, ~74%), Chromium ($\text{Cr}$, ~18%), Nickel ($\text{Ni}$, ~8%), and a trace of Carbon ($\text{C}$, ~0.1%).

Why preferred over pure iron? Pure iron is very soft when hot, stretches easily, and rusts rapidly when exposed to moist air. Stainless steel, on the other hand, is extremely hard, strong, does not rust at all, and is highly resistant to acidic foods. This makes it ideal for making durable household utensils, cutlery, and surgical instruments.

1. Balanced chemical equation: $$\text{Fe}_2\text{O}_3(s) + 2\text{Al}(s) \rightarrow 2\text{Fe}(l) + \text{Al}_2\text{O}_3(s) + \text{Heat}$$
   • Reducing Agent: Aluminium ($\text{Al}$), which gains oxygen to form $\text{Al}_2\text{O}_3$.
   • Oxidizing Agent: Iron(III) oxide ($\text{Fe}_2\text{O}_3$), which loses oxygen.
2. Why iron is molten (liquid): The Thermit reaction is highly exothermic, liberating a massive amount of heat energy (~850 kJ/mol). The temperature generated in the mixture easily exceeds 2500°C, which is far above the melting point of pure iron (1538°C). Therefore, the iron produced melts immediately and flows as a liquid.
3. Why aluminium is used instead of carbon: Aluminium lies higher than iron in the reactivity series and has a much higher chemical affinity for oxygen than carbon does. At temperature limits, aluminium acts as a far more powerful reducing agent, forcing the displacement of iron instantly.

1. Electrodes:
   • Anode (+): Made of the crude impure copper block.
   • Cathode (-): Made of the thin pure copper strip.
   • Reason: Oxidation occurs at the anode, causing copper and more reactive metal impurities to dissolve. Reduction occurs at the cathode, where copper ions gain electrons and deposit as pure metal.
2. Behavior of impurities:
   1. Zinc and Iron (more reactive than Cu): They are oxidized into cations ($\text{Zn}^{2+}$, $\text{Fe}^{2+}$) and dissolve in the copper sulphate bath. They remain dissolved as ions because their reduction potentials are much lower than copper's.
   2. Gold and Silver (less reactive than Cu): They cannot be oxidized at the potential used. As the crude copper anode dissolves, these noble metals simply fall off and collect under the anode as anode mud.
3. Ionic equations at electrodes:
   • At the Anode (+) (Oxidation): $$\text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^-$$
   • At the Cathode (-) (Reduction): $$\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$$

1. Why plain iron is not suitable: Plain iron reacts with water and dissolved oxygen, causing rapid corrosion (rusting). Rust flaky layers peel off into the drinking water, making it dirty and reddish-brown. This clogs the pipes and carries toxic rust particles into drinking water.
2. Cathodic/sacrificial protection concept: Galvanised pipes are coated with zinc. Zinc, being more reactive than iron, oxidizes in preference to iron. If the zinc coating gets scratched, zinc continues to release electrons: $$\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$$ The electrons flow to the exposed iron cathode, keeping the iron reduced and rust-free.
3. Why copper is preferred for hot water lines: At high temperatures, zinc reacts slowly with hot water. However, copper lies below hydrogen in the reactivity series and is completely inert towards hot water and steam. Copper does not corrode or react with water at all, making it highly durable for hot water systems.

1. Chemical name and composition: The liquid is Aqua Regia. It is a freshly prepared fuming mixture of concentrated Hydrochloric acid ($\text{HCl}$) and concentrated Nitric acid ($\text{HNO}_3$) in a 3:1 volume ratio.
2. Why mass decreased: Aqua Regia generates highly reactive nascent chlorine ($\text{Cl}$). When the gold crown is dipped, this chlorine reacts with the outer layer of gold, forming soluble gold chloride ($\text{AuCl}_3$), which dissolves into the liquid: $$\text{Au}(s) + 3\text{Cl}(g) \rightarrow \text{AuCl}_3(aq)$$ Although it removes the tarnish and reveals shiny gold, it actually dissolves a small amount of the gold, reducing the crown's mass. The goldsmith was chemically guilty of dissolving gold.
3. Reactivity of gold: Gold is highly unreactive. It does not react with oxygen, water, or dilute acids. It is called a noble metal because of its chemical inertness, which prevents it from corroding or tarnishing.

1. Why hydrogen evolved instead of sodium: Aqueous sodium chloride contains sodium ions ($\text{Na}^+$), chloride ions ($\text{Cl}^-$), and water molecules which dissociate to yield hydrogen ($\text{H}^+$) and hydroxide ($\text{OH}^-$) ions. In the reactivity series, hydrogen is much lower (less reactive, more easily reduced) than sodium. At the cathode, hydrogen ions gain electrons far more easily than sodium ions do: $$2\text{H}^+(aq) + 2e^- \rightarrow \text{H}_2(g)\uparrow$$ Thus, hydrogen is evolved, and sodium ions remain dissolved.
2. Equations for electrolysis of molten NaCl:
   • At the Cathode (-) (Reduction): $$\text{Na}^+ + e^- \rightarrow \text{Na}(l)$$
   • At the Anode (+) (Oxidation): $$2\text{Cl}^- \rightarrow \text{Cl}_2(g)\uparrow + 2e^-$$
3. Why highly reactive metals require molten reduction: Highly reactive metals have an extremely high affinity for electrons in water, meaning water molecules/hydrogen ions will always be reduced first in aqueous solutions. To prevent water from interfering, we must eliminate water and reduce their molten salts electrolytically.

1. Most reactive metal: Metal B is the most reactive.
   Justification: Metal B is the only metal that can displace iron from iron(II) sulphate solution. Since iron is more reactive than copper, zinc, and silver, a metal capable of displacing iron must be highly reactive (only zinc lies higher, and B did not displace zinc, so B lies between Zn and Fe).
2. Observation if B is added to Copper(II) sulphate: Metal B will undergo a displacement reaction.
   Explanation: Since metal B is more reactive than iron, and iron is more reactive than copper, metal B is significantly more reactive than copper. It will displace copper from its copper sulphate solution:
   • The blue colour of the copper sulphate solution will fade or disappear.
   • A reddish-brown deposit of metallic copper will form on metal B.
3. Decreasing order of reactivity: $$\text{B} > \text{A} > \text{C} > \text{D}$$ Explanation: - B displaces Fe (most reactive). - A does not displace Fe but displaces Cu. - C does not displace Fe or Cu but displaces Ag. - D does not displace any metal (least reactive).

1. Identification:
   • Metal 'X' is Zinc ($\text{Zn}$).
   • White powder 'Y' is Zinc oxide ($\text{ZnO}$), which is amphoteric.
   • Gas 'Z' is Hydrogen gas ($\text{H}_2$).
2. Balanced equations for reactions:
   1. Reaction of 'X' with oxygen: $$2\text{Zn}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{ZnO}(s)$$
   2. Reaction of amphoteric 'Y' with HCl: $$\text{ZnO}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2\text{O}(l)$$
   3. Reaction of amphoteric 'Y' with NaOH: $$\text{ZnO}(s) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{ZnO}_2(aq) + \text{H}_2\text{O}(l)$$ (Sodium zincate is formed).
3. Reaction of 'X' (Zinc) with steam: $$\text{Zn}(s) + \text{H}_2\text{O}(g) \text{ (Steam)} \rightarrow \text{ZnO}(s) + \text{H}_2(g)\uparrow$$ This yields zinc oxide 'Y' and hydrogen gas 'Z'.

This oxide layer acts as a physical shield, preventing further oxygen or moisture from penetrating and reacting with the metal underneath.
Anodising: In industry, this protective oxide layer is artificially thickened by a process called anodising (electrolysing aluminium in sulphuric acid). Because of this stable, non-reactive oxide barrier, aluminium does not corrode or react with food, making it safe and effective for wrapping food and medicines.

Name of the coating: Basic copper carbonate.
Chemical Formula: $\text{CuCO}_3 \cdot \text{Cu(OH)}_2$.
Chemical Equation: $$2\text{Cu}(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g) + \text{O}_2(g) \rightarrow \text{CuCO}_3 \cdot \text{Cu(OH)}_2(s)$$

1. Concentration: The ore is crushed and enriched to remove gangue impurities.
2. Calcination: The carbonate ore is heated strongly in the *absence of air* to decompose it into the metal oxide and carbon dioxide: $$\text{MCO}_3(s) \xrightarrow{\Delta} \text{MO}(s) + \text{CO}_2(g)$$
3. Reduction: The metal oxide ($\text{MO}$) is heated with a suitable reducing agent like carbon (coke) to obtain crude metal 'M': $$\text{MO}(s) + \text{C}(s) \xrightarrow{\Delta} \text{M}(s) + \text{CO}(g)$$
4. Refining: The crude metal is refined (e.g., via electrolytic refining) to yield highly pure metal 'M'.

How it is made hard: It is alloyed with a small amount of harder metals like copper ($\text{Cu}$) or silver ($\text{Ag}$).
What is 22 carat gold? 22 carat gold means a mixture containing 22 parts of pure gold and 2 parts of copper or silver, representing a gold purity of approximately 91.6%.

• Diamond: In diamond, each carbon atom is covalently bonded to four other carbon atoms in a rigid 3D tetrahedral network. All valence electrons are tightly locked in localized single covalent bonds, leaving no free electrons. Thus, diamond is an electrical insulator.
• Graphite: In graphite, each carbon atom is bonded to only three other carbon atoms in a hexagonal planar layer. This leaves one valence electron per carbon atom completely free and delocalised. These mobile electrons can move easily along the layers, making graphite an excellent conductor.

Why used as electrodes: Graphite conducts electricity efficiently, has a high melting point, and is chemically inert and cheap, making it ideal for inert electrodes.

Balanced chemical equation: $$3\text{Fe}(s) + 2\text{O}_2(g) \xrightarrow{\Delta} \text{Fe}_3\text{O}_4(s)$$ Product Name: Iron(II,III) oxide or magnetic iron oxide ($\text{Fe}_3\text{O}_4$).

1. Oxygen Molecule ($\text{O}_2$): Oxygen (atomic number 8, configuration 2, 6) has 6 valence electrons. Two oxygen atoms share 2 pairs of electrons to complete their octets, forming a covalent double bond: $$\text{:}\ddot{\text{O}}\text{:} + \text{:}\ddot{\text{O}}\text{:} \rightarrow \text{:}\ddot{\text{O}}\text{::}\ddot{\text{O}}\text{:} \rightarrow \text{O}=\text{O}$$
2. Nitrogen Molecule ($\text{N}_2$): Nitrogen (atomic number 7, configuration 2, 5) has 5 valence electrons. Two nitrogen atoms share 3 pairs of electrons to complete their octets, forming a covalent triple bond: $$\text{:}\dot{\text{N}}\text{:} + \text{:}\dot{\text{N}}\text{:} \rightarrow \text{:N:::N:} \rightarrow \text{N}\equiv\text{N}$$

Non-metals behavior: Non-metals are electronegative elements. They have high electron affinities and tend to gain electrons rather than lose them. Because non-metals cannot donate electrons, they cannot reduce $\text{H}^+$ ions to hydrogen gas. Hence, non-metals do not displace hydrogen from acids.

The description identifies 'X' as an ionic compound (e.g., sodium chloride).

1. High Solubility & Conductivity: Ionic compounds are highly polar and dissolve easily in polar water molecules. In solution, the polar water molecules break the lattice, releasing mobile hydrated ions which conduct electricity efficiently.
2. Brittleness under Hammer: Ionic crystals consist of alternating layers of closely packed cations and anions. When struck with a hammer, the impact forces the layers of the crystal lattice to shift slightly. This shifting brings ions of like charges directly opposite each other (e.g., cation opposite cation, anion opposite anion). The resulting sudden, massive electrostatic repulsion shatters the crystal along cleavage planes, showing brittleness.

1. Red Rust (Plain unprotected iron): Highly ineffective. Rust is a hydrated iron oxide ($\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$) that forms a flaky, porous coating. It does not stick to the metal surface, continuously peeling off and exposing fresh iron underneath. Over time, it completely eats away the structural iron, leading to catastrophic failure.
2. Grey Zinc Paint (Galvanic paint coating): Highly effective and durable. This works in two ways. First, it provides a physical barrier. Second, even if scratched, zinc is more reactive than iron and acts as a sacrificial anode, corroding in preference to the iron bridge structure. It provides long-lasting cathodic protection.
3. Brass Bolts (Alloy of Cu and Zn): Extremely effective for small structural joints. Brass is highly resistant to corrosion because the homogeneous mixing of zinc and copper alters the atomic lattice, making it chemically inert to atmospheric oxygen and water. It maintains its structural strength indefinitely, though it is too expensive to build the entire bridge from.