A Short Revision in Physical Chemistry for Undergraduate

 

1. Introduction to Physical Chemistry 

Physical chemistry deals with the quantitative and theoretical aspects of chemistry, including:

• Thermodynamics (energy changes in reactions).

• Kinetics (reaction rates and mechanisms).

• Quantum Mechanics (molecular structure and bonding).

• Electrochemistry (redox reactions, batteries, corrosion).

• Spectroscopy (IR, UV, NMR, mass spectrometry).

• Statistical Mechanics (molecular behavior in gases and solutions).

2. Thermodynamics (Energy & Spontaneity of Reactions)

Laws of Thermodynamics

First Law: Energy cannot be created or destroyed, only transferred.

$$Delta U = q + w$$

(Change in internal energy = heat + work done on the system)

Second Law: Entropy (disorder) always increases in a spontaneous process.

$$\mathrm{\Delta }{S}_{\text{universe}}>0$$

Third Law: A perfect crystal at absolute zero (0 K) has zero entropy.

Enthalpy ($Delta H$) – Heat of Reaction

• Exothermic:

• $Delta H$ (heat released, e.g., combustion).

• Endothermic: $Delta H > 0$ (heat absorbed, e.g., melting ice).

Hess’s Law: The total enthalpy change in a reaction is independent of the pathway.

Gibbs Free Energy ($Delta G$) – Spontaneity of Reactions

$$Delta G = Delta H - Delta S$$

• $Delta G$ → Spontaneous reaction

• $Delta G > 0$ → Non-spontaneous reaction

• $Delta G = 0$ → System at equilibrium

Example: Why does ice melt at room temperature?

• $Delta H > 0$ (heat absorbed).

• $Delta S > 0$ (disorder increases).

• $Delta G < 0$ at T > 0°C, so melting is spontaneous.

3. Chemical Kinetics (Reaction Rates & Mechanisms)

Rate Laws

For a reaction A + B → C, the rate equation is:

$$k[A]^m[B]^n$$

where k is the rate constant and m, n are reaction orders.

Factors Affecting Reaction Rates

• Concentration: More reactants → Faster reaction.

• Temperature: Higher

• $T$ → Faster reaction (Arrhenius equation).

• Catalysts: Lower activation energy ($E_a$).

Arrhenius Equation (Temperature Dependence of Rates)

$$\mathrm{<br>}$$

Example: Why does food spoil faster in summer?

• Higher $T$ → Larger $k$ → Faster reaction rate.

4. Chemical Equilibrium

When forward and backward reaction rates are equal:

• $K > 1$ → Products favored.

• $K < 1$ → Reactants favored.

Le Châtelier’s Principle: If a system at equilibrium is disturbed, it shifts to counteract the change.

5. Electrochemistry (Redox Reactions & Batteries)

Galvanic Cells (Batteries)

• Spontaneous redox reaction ($\Delta G < 0$) generates electricity.

• Example: Zn-Cu Daniell cell

  $$Zn(s)+C{u}^{2+}(aq)\to Z{n}^{2+}(aq)+Cu(s)$$

Nernst Equation (Cell potential at non-standard conditions)

$${}^{}$$

Electrolytic Cells (Electrolysis)

• Non-spontaneous reaction ($\Delta G > 0$), requiring external energy (e.g., electrolysis of water into H₂ and O₂).

6. Quantum Chemistry & Atomic Structure

Wave-Particle Duality (de Broglie’s Hypothesis)

$$\lambda =\frac{h}{mv}$$

Electrons behave as both particles and waves, leading to orbitals instead of orbits.

Schrödinger Equation & Atomic Orbitals

• s, p, d, f orbitals describe electron probability clouds.

• Pauli Exclusion Principle: No two electrons have the same set of quantum numbers.

7. Spectroscopy (Analytical Chemistry)

Infrared (IR) Spectroscopy – Functional Groups

• C=O (carbonyl group) → 1700 cm⁻¹

• O-H (alcohols, acids) → 3200-3600 cm⁻¹

Nuclear Magnetic Resonance (NMR) Spectroscopy,  Structure Determination

• ¹H NMR: Identifies hydrogen environments.

• ¹³C NMR: Identifies carbon environments.

Mass Spectrometry (MS) – Molecular Mass

• Molecular ion peak ($M^+$) → Exact mass of compound.

That wraps up our Physical Chemistry revision for undergraduate !