Physical Chemistry Master physical chemistry with our chemistry tuition Singapore program, featuring clear explanations and exam-focused practice. covers thermodynamics, kinetics, and equilibrium. This guide explains the key concepts.
Atomic Structure
Key Concepts:
- Quantum numbers
- – Electronic configuration
- – Orbital shapes
- – Periodic trends
Electronic Configuration:
- Aufbau principle
- – Pauli exclusion principle
- – Hund’s rule
Example: Fe (Z=26)
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
or [Ar] 3d⁶ 4s²
Chemical Bonding
Types of Bonding:
- Ionic
- 2. Covalent
- 3. Metallic
Ionic Bonding:
- Electrostatic attraction
- – Lattice structure
- – High melting points
Covalent Bonding:
- Shared electrons
- – Sigma and pi bonds
- – VSEPR theory
VSEPR Shapes:
- Linear (2 pairs)
- – Trigonal planar (3 pairs)
- – Tetrahedral (4 pairs)
- – Trigonal bipyramidal (5 pairs)
- – Octahedral (6 pairs)
Intermolecular Forces:
- Van der Waals
- – Dipole-dipole
- – Hydrogen bonding
Thermodynamics
Enthalpy (ΔH):
- Heat change at constant pressure
- – ΔH = H(products) – H(reactants)
Types:
- ΔHf: formation
- – ΔHc: combustion
- – ΔHneut: neutralisation
- – ΔHhyd: hydration
- – ΔHsol: solution
Hess’s Law:
Total enthalpy change is independent of pathway.
Application:
- Calculate unknown ΔH
- – Use enthalpy cycles
Entropy (ΔS):
- Measure of disorder
- – ΔS = S(products) – S(reactants)
Factors affecting ΔS:
- State changes
- – Number of particles
- – Complexity
Gibbs Free Energy (ΔG):
- ΔG = ΔH – TΔS
- – Predicts spontaneity
ΔG < 0: spontaneous
ΔG > 0: non-spontaneous
ΔG = 0: equilibrium
Chemical Equilibrium
Dynamic Equilibrium:
- Forward and backward rates equal
- – Concentrations constant
- – Closed system
Equilibrium Constant (Kc):
aA + bB ⇌ cC + dD
Kc = [C]^c [D]^d / [A]^a [B]^b
Factors Affecting Equilibrium:
- Concentration
- 2. Pressure (gases)
- 3. Temperature
Le Chatelier’s Principle:
System opposes changes to restore equilibrium.
Effect of Catalyst:
- Increases rate
- – No effect on position
- – Lowers activation energy
Reaction Kinetics
Rate of Reaction:
- Change in concentration per unit time
- – Affected by: concentration, temperature, catalysts
Rate Equation:
Rate = k[A]^m[B]^n
Where:
- k = rate constant
- – m, n = orders of reaction
- – Overall order = m + n
Determining Orders:
- From experimental data
- – Initial rates method
- – Concentration-time graphs
Half-life (t½):
- Time for concentration to halve
- – First order: t½ = ln2/k
Arrhenius Equation:
k = Ae^(-Ea/RT)
Where:
- Ea = activation energy
- – A = pre-exponential factor
- – R = gas constant
- – T = temperature
From a graph of ln k vs 1/T:
Gradient = -Ea/R
Acid-Base Equilibria
Bronsted-Lowry Theory:
- Acid: proton donor
- – Base: proton acceptor
Strong vs Weak:
- Strong: complete dissociation
- – Weak: partial dissociation
pH:
pH = -log[H+]
Kw:
Kw = [H+][OH-] = 10^-14 at 25°C
Ka and Kb:
- Ka: acid dissociation constant
- – Kb: base dissociation constant
- – Ka × Kb = Kw
Buffers:
- Resist pH changes
- – Contain weak acid + conjugate base
- – Or weak base + conjugate acid
Electrochemistry
Redox Reactions:
- Oxidation: loss of electrons
- – Reduction: gain of electrons
Electrochemical Cells:
- Anode: oxidation
- – Cathode: reduction
Cell Potential:
E°cell = E°cathode – E°anode
Spontaneous if E°cell > 0
Electrolysis:
- Uses electrical energy
- – Non-spontaneous reactions
Common Mistakes
- Wrong units
- 2. Sign errors
- 3. Not balancing equations
- 4. Misapplying formulas
- 5. Wrong equilibrium expressions
Study Tips
- Understand concepts
- 2. Memorise key formulas
- 3. Practise calculations
- 4. Draw diagrams
- 5. Review regularly
How Ace Scorers Helps
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