METALS AND THE REACTIVITY SERIES

Quiz 1:
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🧪 METALS & REACTIVITY SERIES: Complete Theory & Formulas

🔬 I. FUNDAMENTAL THEORIES OF METALS

1.1 Electronic Theory of Metallic Bonding

Sea of Electrons Model: Metals consist of a lattice of positive metal ions surrounded by a "sea" of delocalized electrons. This explains conductivity, malleability, and ductility.

Metal Atom → Metal Ion⁺ + e⁻

M(s) → M⁺(aq) + e⁻ (general form)

1.2 Density Theory & Formula

Density is a fundamental physical property that measures mass per unit volume.

Density = Mass ÷ Volume

ρ = m/V (g/cm³ or kg/m³)

High Density Examples

• Gold: 19.3 g/cm³
• Lead: 11.4 g/cm³
• Copper: 8.9 g/cm³

Low Density Examples

• Sodium: 0.97 g/cm³
• Magnesium: 1.7 g/cm³
• Aluminum: 2.7 g/cm³

⚡ II. REACTIVITY SERIES THEORY

2.1 Electron Loss Tendency Theory

The reactivity series is based on the tendency of metals to lose electrons and form positive ions. More reactive metals lose electrons more readily.

Complete Reactivity Series

K (Potassium) - Most Reactive

Na (Sodium)

Ca (Calcium)

Mg (Magnesium)

Al (Aluminum)

C (Carbon) - Reference

Zn (Zinc)

Fe (Iron)

Pb (Lead)

H (Hydrogen) - Reference

Cu (Copper)

Ag (Silver)

Au (Gold) - Least Reactive

2.2 Displacement Reaction Theory

A more reactive metal can displace a less reactive metal from its compound because it has a stronger tendency to form positive ions.

General Form: A + BC → AC + B

Where A is more reactive than B

🧪 III. CHEMICAL REACTION FORMULAS

3.1 Reactions with Water

Potassium + Water:
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Sodium + Water:
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Calcium + Water:
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Magnesium + Steam:
Mg(s) + H₂O(g) → MgO(s) + H₂(g)

3.2 Reactions with Acids

General Formula:
Metal + Acid → Salt + Hydrogen

Magnesium + Hydrochloric Acid:
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

Zinc + Sulfuric Acid:
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

Iron + Hydrochloric Acid:
Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

3.3 Reactions with Oxygen

General Formula:
Metal + Oxygen → Metal Oxide

Magnesium + Oxygen:
2Mg(s) + O₂(g) → 2MgO(s)

Iron + Oxygen (Rusting):
4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s)

🏭 IV. EXTRACTION THEORIES & METHODS

4.1 Reduction Theory

Metals below carbon in the reactivity series can be extracted by reduction with carbon because carbon is more reactive than these metals.

Reactivity Level	Extraction Method	Examples
Above Carbon	Electrolysis	K, Na, Ca, Mg, Al
Below Carbon	Reduction with Carbon	Zn, Fe, Pb, Cu
Very Unreactive	Found Naturally	Ag, Au

4.2 Iron Extraction (Blast Furnace)

Four-stage reduction process in blast furnace:

Stage 1 - Heat Generation:
C(s) + O₂(g) → CO₂(g)

Stage 2 - Reducing Agent Formation:
C(s) + CO₂(g) → 2CO(g)

Stage 3 - Iron Reduction:
Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)

Stage 4 - Impurity Removal:
CaCO₃(s) → CaO(s) + CO₂(g)
CaO(s) + SiO₂(s) → CaSiO₃(l)

4.3 Aluminum Extraction (Electrolysis)

Two-stage electrolytic process:

Overall Reaction:
2Al₂O₃(l) → 4Al(l) + 3O₂(g)

At Cathode (Reduction):
Al³⁺(l) + 3e⁻ → Al(l)

At Anode (Oxidation):
2O²⁻(l) → O₂(g) + 4e⁻

Anode Consumption:
C(s) + O₂(g) → CO₂(g)

🛡️ V. CORROSION THEORY

5.1 Rusting Theory

Rusting is an oxidation process requiring both oxygen and water. It follows the electrochemical theory of corrosion.

Fe(s) → Fe³⁺(aq) + 3e⁻ (Oxidation at anode)

O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) (Reduction at cathode)

5.2 Sacrificial Protection Theory

Based on the reactivity series - a more reactive metal will oxidize preferentially, protecting the less reactive metal.

Zinc Protection of Iron:
Zn(s) → Zn²⁺(aq) + 2e⁻ (Zinc oxidizes first)

Electron Flow:
Electrons flow from Zn to Fe, preventing Fe oxidation

⚗️ VI. ALLOY THEORY

6.1 Structural Theory of Alloys

Alloys are harder than pure metals due to size difference of atoms disrupting the regular crystal lattice structure.

⚡ VII. ELECTROCHEMICAL THEORIES

7.1 Simple Cell Theory

Based on difference in electrode potential between two metals in the reactivity series.

Cell Voltage ∝ Difference in Reactivity

More Reactive Metal = Negative Terminal (Anode)

Less Reactive Metal = Positive Terminal (Cathode)

7.2 Thermite Reaction Theory

Highly exothermic displacement reaction based on large difference in reactivity between aluminum and iron.

Thermite Reaction:
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat

Enthalpy Change:
ΔH = -850 kJ/mol (highly exothermic)

🌡️ VIII. THERMAL DECOMPOSITION THEORY

8.1 Stability-Reactivity Relationship

The lower the metal in reactivity series, the more easily its compounds decompose when heated.

Metal Carbonates (except Na₂CO₃, K₂CO₃):
MCO₃(s) → MO(s) + CO₂(g)

Metal Hydroxides (except NaOH, KOH):
M(OH)₂(s) → MO(s) + H₂O(g)

Metal Nitrates (most metals):
2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g)

Sodium/Potassium Nitrates:
2MNO₃(s) → 2MNO₂(s) + O₂(g)

🎯 Key Applications & Modern Uses

• Aerospace Industry: Aluminum alloys for lightweight, strong aircraft structures
• Electronics: Copper for conductivity, gold for corrosion resistance in circuits
• Construction: Steel (iron alloy) with sacrificial zinc coating (galvanizing)
• Transportation: Various alloys optimized for specific strength-to-weight ratios
• Medical: Titanium alloys for biocompatible implants

💡 Remember: The Universal Principles

1. Electron Theory: All metallic properties stem from electron mobility
2. Reactivity Hierarchy: Determines all chemical behavior patterns
3. Thermodynamic Stability: More reactive = more stable compounds
4. Extraction Economics: Method difficulty correlates with reactivity
5. Practical Applications: Properties determine technological uses

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