Chemistry 8A

Quiz 1:
 Fullscreen Mode

Complete Chemistry Theory: Bonding, Reaction Rates, and Equilibrium

Section 1: Ionic and Covalent Bonding

Atomic Structure

Atoms consist of three fundamental particles:

• Protons: Positively charged particles in the nucleus
• Neutrons: Neutral particles in the nucleus
• Electrons: Negatively charged particles orbiting the nucleus

Ion Formation

Cation Formation: Atom → Cation + e⁻
Anion Formation: Atom + e⁻ → Anion

Na⁺
Cation

+

Cl⁻
Anion

→

NaCl
Ionic Compound

The Octet Rule

Stable Configuration: 8 electrons in outer shell
(Exception: Hydrogen needs only 2 electrons)

Types of Chemical Bonds

Ionic Bonds

Formation: Metal + Non-metal → Ionic Compound
Mechanism: Complete electron transfer
Example: Na + Cl → Na⁺Cl⁻ (NaCl)

Ionic Bond Energy ∝ (q₁ × q₂) / r
Where: q₁, q₂ = charges, r = distance between ions

Covalent Bonds

Formation: Non-metal + Non-metal → Covalent Compound
Mechanism: Electron sharing
Example: C + 2O → CO₂

Properties Comparison

Ionic Compounds:

• High melting points (1000°C - 3000°C)
• Conduct electricity when dissolved
• Hard but brittle
• Form crystalline structures

Covalent Compounds:

• Low melting points (< 200°C)
• Poor electrical conductors
• Can be gases, liquids, or soft solids
• Form molecular structures

Section 2: Rate of Reaction

Definition and Measurement

Rate of Reaction = Δ[Product] / Δt = -Δ[Reactant] / Δt
Units: mol dm⁻³ s⁻¹ or g s⁻¹

Collision Theory

For a reaction to occur:

1. Particles must collide
2. Collision must have sufficient energy (≥ activation energy)
3. Particles must have correct orientation

Activation Energy

Ea = Minimum energy required for reaction
Rate ∝ e^(-Ea/RT)

Factors Affecting Reaction Rate

1. Concentration

Rate ∝ [A]ᵐ[B]ⁿ
Where m, n = reaction orders

2. Temperature

Arrhenius Equation: k = Ae^(-Ea/RT)
Where: k = rate constant, A = frequency factor, R = gas constant, T = temperature

3. Surface Area

Effect: Increased surface area → More collision sites → Higher rate
Example: Powdered marble reacts faster with HCl than marble chips

4. Catalysts

Catalyst: Lowers Ea without being consumed
Alternative pathway: Reactants → Intermediate → Products

Types of Reactions

Exothermic Reactions:

ΔH < 0 (Energy released)

Example: Combustion, Neutralization

Endothermic Reactions:

ΔH > 0 (Energy absorbed)

Example: Photosynthesis, Thermal decomposition

Section 3: Reversible Reactions and Equilibrium

Reversible Reactions

General Form: aA + bB ⇌ cC + dD
Forward Rate = Reverse Rate (at equilibrium)

Dynamic Equilibrium

Characteristics:

• Closed system required
• Constant macroscopic properties
• Continuous molecular activity
• Forward rate = Reverse rate

Equilibrium Constant

Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
(For reaction: aA + bB ⇌ cC + dD)

Le Chatelier's Principle

Principle: If a system at equilibrium is disturbed, it shifts to counteract the disturbance.

Effects of Changes:

• Concentration: Adding reactant shifts right; adding product shifts left
• Temperature: Increasing T favors endothermic direction
• Pressure: Increasing P favors side with fewer gas molecules
• Catalyst: No effect on position, only rate to reach equilibrium

Examples of Reversible Reactions

Copper Sulfate Hydration

CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(g)
Blue (hydrated) ⇌ White (anhydrous) + Water

Cobalt Chloride System

CoCl₂ + 6H₂O ⇌ CoCl₂·6H₂O
Blue + Water ⇌ Pink (hydrated)

Position of Equilibrium

Left-favored: Kc < 1 (More reactants)
Right-favored: Kc > 1 (More products)
Balanced: Kc ≈ 1 (Similar amounts)

Irreversible vs Reversible Reactions

Irreversible Examples:

• Combustion: C + O₂ → CO₂
• Precipitation reactions
• Acid-base neutralization (complete)

Reversible Examples:

• Phase changes: H₂O(l) ⇌ H₂O(g)
• Weak acid dissociation: CH₃COOH ⇌ CH₃COO⁻ + H⁺
• Ester formation: Acid + Alcohol ⇌ Ester + Water

Key Takeaways

Understanding chemical bonding, reaction rates, and equilibrium provides the foundation for predicting and controlling chemical processes in both laboratory and industrial applications.