1. Thermodynamic Terms
Thermodynamics deals with energy changes in chemical and physical processes. Key definitions include:
- System: The part of the universe under observation (e.g., contents of a flask).
- Surroundings: Everything else in the universe outside the system.
- Boundary: The real or imaginary surface separating the system from the surroundings.
Types of Systems
| System Type | Exchange of Matter | Exchange of Energy | Analogy |
|---|---|---|---|
| Open (e.g., open beaker) | Yes | Yes | A door standing wide open. |
| Closed (e.g., sealed flask) | No | Yes | A closed window (you see light/heat but can't pass through). |
| Isolated (e.g., thermos) | No | No | A thick, vault-like wall. |
Intensive vs Extensive Shortcut
- Extensive Properties: Depend on amount. (Mass, Volume, U, H, S, G). Think: "eXpands with size".
- Intensive Properties: Independent of amount. (Temp, Pressure, Density, pH, Cell Potential). Think: "INdependent".
Gold Tip: Specific properties (e.g., Specific Heat) and Molar properties (e.g., Molar Volume) are always Intensive.
State and Path Functions
State Functions: depend only on initial and final states. (P, V, T, U, H, S, G). Most functions in thermo are state functions!
Path Functions: depend on the mechanism. Only two: Heat (q) and Work (w).
2. Calorimetry (ΔU and ΔH)
Calorimetry measures heat flow. For IAT, remember the distinction between Constant Volume and Constant Pressure.
| Technique | Condition | Heat (q) equals... | Concept |
|---|---|---|---|
| Bomb Calorimetry | ΔV = 0 | qv = ΔU | Used for combustion. No expansion work. |
| Coffee Cup | ΔP = 0 | qp = ΔH | Used for solutions at atmospheric pressure. |
IAT Logic: Cp vs Cv
Always remember: Cp > Cv. Why? At constant pressure, part of the heat supplied is used for expansion work (PΔV).
For ideal gases: Cp - Cv = R.
3. Enthalpy Change (ΔH)
The relation between ΔH and ΔU is one of the most tested IAT topics.
Lattice Enthalpy & Born-Haber Cycle
Lattice Enthalpy: Energy required to break 1 mole of ionic solid into gaseous ions. It depends on Charge (directly) and Size (inversely).
The Born-Haber Cycle uses Hess's Law to find lattice enthalpy by summing sublimation, I.E., bond energy, and E.A.
4. Enthalpies of Reaction
| Type | Symbol | Critical Point |
|---|---|---|
| Standard Formation | ΔfH° | Formation of 1 mole from elements in stable state. (Stable elements = 0). |
| Neutralization | ΔneutH° | Strong Acid + Strong Base = -57.1 kJ/mol. Weak acids/bases give less heat. |
| Combustion | ΔcH° | Always Exothermic (negative). |
5. Spontaneity and Entropy
Entropy (S): A measure of randomness. Second Law of Thermodynamics states: For a spontaneous process, ΔStotal > 0.
Entropy Change Rules:
- Gas > Liquid > Solid (Highest entropy in gases).
- Entropy increases with Temperature and Volume.
- Entropy increases if moles of gas increase during a reaction.
6. Gibbs Energy & Equilibrium
The absolute criteria for spontaneity at constant T and P is ΔG < 0.
Gibbs and Equilibrium Constant (K)
Relation: ΔG = ΔG° + RT ln Q
At Equilibrium (ΔG = 0, Q = K): ΔG° = -RT ln K = -2.303 RT log K
Quick Revision Flashcards
Hover over a card to reveal the definition/formula.
Hess's Law
Total enthalpy change is independent of the path taken. (State Function property).
Third Law
The entropy of a pure crystalline substance is ZERO at absolute zero (0 K).
Adiabatic
Process with NO heat exchange (q = 0). For adiabatic change: ΔU = w.
7. Common Mistakes & IAT Prep
- Unit Trap: Always convert ΔS (J/K) to kJ/K by dividing by 1000 before using the Gibbs equation.
- Free Expansion: Expansion into vacuum means Pext = 0, thus w = 0.
- Exothermic vs Spontaneous: Not all exothermic reactions are spontaneous, and not all spontaneous reactions are exothermic. ΔG is the final judge!
9. Practice Mock Test
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Thermodynamics