Thermodynamics – Study Notes
(Program: Military Nursing Service – B.Sc. Chemistry)
1. Simple Explanation
Thermodynamics is the branch of chemistry that deals with the relationships between heat, work, energy, and the physical properties of matter. In a nutshell:
| Concept | What it means | Why it matters in nursing? |
|---|---|---|
| Energy | Anything that can do work or be stored. | Body heat, metabolism, drug action. |
| Heat (q) | Energy transfer due to temperature difference. | Fever, hypothermia, wound healing. |
| Work (w) | Energy transfer by a force acting through a distance. | Mechanical ventilation, movement. |
| State Functions | Properties that depend only on the state of a system (T, P, V, etc.). | Enthalpy, entropy, Gibbs free energy. |
| Reversible vs. Irreversible | Ideal vs. real processes. | Drug administration (slow release vs. injection). |
2. Key Points
| Topic | Key Equations | Key Take‑Aways |
|---|---|---|
| First Law (Energy Conservation) | ΔU = q + w | Energy in a system changes only by heat added or work done. |
| Enthalpy (H) | H = U + PV | Useful for reactions at constant pressure (e.g., body temperature). |
| Heat Capacity (C) | C = q/ΔT | Determines how much heat a substance needs to change temperature. |
| Second Law (Entropy, ΔS) | ΔS = q_rev/T | Spontaneous processes increase total entropy (system + surroundings). |
| Gibbs Free Energy (ΔG) | ΔG = ΔH – TΔS | Predicts spontaneity at constant T & P. |
| Equilibrium Constant (K) | ΔG° = –RT ln K | Connects thermodynamics to chemical equilibrium. |
| Heat of Reaction (ΔH_rxn) | ΔH_rxn = Σ ΔH_f (products) – Σ ΔH_f (reactants) | Exothermic vs. endothermic reactions. |
| Phase Changes | ΔH_fusion, ΔH_vaporization | Important for body temperature regulation. |
| Temperature Dependence of K | ln K = –ΔH°/RT + ΔS°/R | Van ’t Hoff equation. |
Quick glossary
- Δ (Delta) = change in a property.
- q = heat (positive = absorbed, negative = released).
- w = work (positive = done on system, negative = done by system).
- U = internal energy.
- PV = pressure × volume.
- R = ideal gas constant (8.314 J mol⁻¹ K⁻¹).
- T = absolute temperature (K).
3. Worked Example
Problem
Calculate the Gibbs free energy change (ΔG) for the reaction
[ \text{NH}_4^+ + \text{OH}^- \rightarrow \text{NH}_3 + \text{H}_2\text{O} ]
at 25 °C.
Given:
- ΔH° = –25 kJ mol⁻¹
- ΔS° = –50 J mol⁻¹ K⁻¹
Step‑by‑Step
-
Convert ΔS° to kJ (since ΔH° is in kJ):
[ \Delta S^\circ = -50\ \text{J mol}^{-1}\text{K}^{-1} = -0.050\ \text{kJ mol}^{-1}\text{K}^{-1} ]
-
Temperature in Kelvin:
[ T = 25,^\circ\text{C} = 298,\text{K} ]
-
Use ΔG° = ΔH° – TΔS°:
[ \Delta G^\circ = (-25\ \text{kJ mol}^{-1}) - (298\ \text{K})(-0.050\ \text{kJ mol}^{-1}\text{K}^{-1}) ]
-
Compute the second term:
[ 298 \times -0.050 = -14.9\ \text{kJ mol}^{-1} ]
-
Add the terms:
[ \Delta G^\circ = -25 - (-14.9) = -25 + 14.9 = -10.1\ \text{kJ mol}^{-1} ]
Result
[ \boxed{\Delta G^\circ = -10.1\ \text{kJ mol}^{-1}} ]
Interpretation – The negative ΔG° indicates the reaction is spontaneous at 25 °C under standard conditions.
4. Quick Revision Summary
| Concept | Formula | Sign of Δ? | What it tells you |
|---|---|---|---|
| First Law | ΔU = q + w | ΔU can be positive or negative | Energy balance |
| Enthalpy | H = U + PV | ΔH can be + (endothermic) or – (exothermic) | Heat change at constant P |
| Heat Capacity | C = q/ΔT | N/A | Amount of heat to change 1 K |
| Entropy | ΔS = q_rev/T | ΔS > 0 for spontaneous | Disorder increase |
| Gibbs Free Energy | ΔG = ΔH – TΔS | ΔG < 0 → spontaneous | Predicts reaction direction |
| Equilibrium | ΔG° = –RT ln K | K > 1 → products favored | Equilibrium position |
| Van ’t Hoff | d(ln K)/dT = ΔH°/RT² | If ΔH° > 0, K ↑ with T | Temperature effect on K |
Quick Tips for Exams
- **Always