Unit 2: Chemical Thermodynamics
Course Code: CHM-DSM-252
Paper Name: Fundamentals of Chemistry - II
1. Review of Laws of Thermodynamics
Thermodynamics is the study of energy, heat, work, and their interconversions in chemical systems.
- First Law: Energy can neither be created nor destroyed, only transformed from one form to another.
- Second Law: In any spontaneous process, the total entropy of the universe always increases.
- Third Law: The entropy of a perfectly crystalline substance at absolute zero (0 K) is exactly zero.
2. Principles of Thermochemistry
Thermochemistry focuses on the heat energy absorbed or released during chemical reactions.
Key Definitions
- System: The specific part of the universe under study.
- Surroundings: Everything outside the system.
- Exothermic Reaction: A reaction that releases heat to the surroundings (ΔH is negative).
- Endothermic Reaction: A reaction that absorbs heat from the surroundings (ΔH is positive).
3. Standard Enthalpies of Formation and Solution
To standardize measurements, thermodynamic properties are reported under "Standard State" conditions.
Standard Enthalpy of Formation (ΔHf°)
The enthalpy change when one mole of a substance is formed from its elements in their most stable states under standard conditions.
Enthalpies of Solution and Dilution
- Integral Enthalpy of Solution: The heat change when a specific amount of solute is dissolved in a solvent to form a solution of a specific concentration.
- Differential Enthalpy of Solution: The heat change when one mole of solute is added to a very large volume of solution so that the concentration remains virtually unchanged.
- Enthalpy of Dilution: The heat change that occurs when more solvent is added to a solution.
4. Bond Energy and Resonance Energy
Thermochemical data allows for the calculation of the energy required to break chemical bonds.
Bond Energy and Bond Dissociation Energy
- Bond Dissociation Energy: The energy required to break one mole of a specific bond in a gaseous molecule.
- Bond Energy: Often used as the average value of bond dissociation energies for similar bonds in different molecules.
Resonance Energy
Resonance energy is the difference between the actual (experimental) enthalpy of formation of a compound and the value calculated for its most stable Lewis structure.
Resonance Energy = ΔH (experimental) - ΔH (calculated)
5. Temperature Dependence: Kirchhoff’s Equation
Kirchhoff’s equation describes how the enthalpy change (ΔH) of a chemical reaction varies with temperature.
ΔH2 - ΔH1 = ΔCp (T2 - T1)
Where ΔCp is the difference in molar heat capacities at constant pressure between products and reactants:
ΔCp = Σ Cp (products) - Σ Cp (reactants).
6. Exam Focus: Tips and FAQs
Exam Tips:
- State Functions: Remember that Enthalpy (H) and Internal Energy (U) are state functions, meaning they depend only on the initial and final states, not the path taken.
- Standard State: In calculations, the standard enthalpy of formation for elements in their stable form (e.g., O2 gas, C graphite) is taken as zero.
- Kirchhoff's Equation: This is a common derivation; ensure you understand the relationship between ΔH and heat capacity.
Frequently Asked Questions
Q: What is the difference between integral and differential enthalpy of solution?
A: Integral enthalpy refers to dissolving a whole amount of solute into a solvent, while differential enthalpy refers to adding a tiny amount of solute to an existing solution.
Q: How is bond energy calculated from thermochemical data?
A: It is calculated using Hess’s Law by comparing the enthalpies of formation of gaseous atoms and the molecule itself.
Q: Define standard state.
A: It is the most stable physical state of a substance at 1 bar pressure and a specified temperature (usually 298.15 K).