Thermodynamics

Learning Objectives

• Understand the mathematical expression of the First Law of Thermodynamics, $U = q + W$, and its application in energy conservation for thermodynamic processes.
• Classify thermodynamic systems as open, closed, or isolated based on their interactions with surroundings in terms of matter and energy exchange.
• Analyze the relationship between enthalpy ($H$) and internal energy ($U$) in chemical reactions, particularly using the equation $H = U + pV$ for processes at constant pressure.
• Apply Hess's Law of Constant Heat Summation to calculate reaction enthalpies from known enthalpies of formation, emphasizing its path-independent nature.
• Calculate bond enthalpies and use them to estimate reaction enthalpies, focusing on the energy required to break specific bonds in gaseous molecules.
• Explain lattice enthalpy as the enthalpy change when one mole of an ionic compound dissociates into its gaseous ions, and utilize the Born-Haber cycle for calculations.
• Determine standard enthalpy changes for various processes such as formation, combustion, fusion, and vaporization under standard conditions.
• Conduct calorimetry experiments to measure heat changes in chemical reactions, utilizing devices like bomb calorimeters for constant volume processes.
• Differentiate between spontaneous and non-spontaneous processes, and understand why the First Law alone cannot predict spontaneity.
• Interpret entropy as a measure of disorder/randomness and apply the Second Law of Thermodynamics to determine the spontaneity of processes.
• Evaluate Gibbs energy changes ($G = H - TS$) to determine spontaneity and equilibrium constants, using the criteria $\Delta G < 0$, $\Delta G = 0$, and $\Delta G > 0$.

Chapter Notes

First Law of Thermodynamics

• Mathematical Expression: $U = q + W$
  • $U$: Change in internal energy
  • $q$: Heat added to the system
  • $W$: Work done by the system
• Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Types of Thermodynamic Systems

• Open System: Exchanges both matter and energy with surroundings.
• Closed System: Exchanges energy but not matter with surroundings.
• Isolated System: No exchange of matter or energy with surroundings.

Enthalpy and Internal Energy

• Relationship: $H = U + pV$
  • $H$: Enthalpy
  • $U$: Internal energy
  • $pV$: Pressure-volume work
• Constant Pressure Process: Enthalpy change equals heat absorbed.

Hess's Law of Constant Heat Summation

• Principle: Total enthalpy change for a reaction is the same regardless of the pathway.
• Application: Allows calculation of reaction enthalpies from known enthalpies of formation.

Bond Enthalpy

• Definition: Energy required to break one mole of a specific type of bond in a gaseous molecule.
• Use: Estimate reaction enthalpies.

Lattice Enthalpy

• Definition: Enthalpy change when one mole of an ionic compound dissociates into its gaseous ions.
• Calculation: Often uses the Born-Haber cycle.

Standard Enthalpy Changes

• Processes: Formation, combustion, fusion, vaporization.
• Conditions: Standard state conditions (1 bar, specified temperature).

Calorimetry

• Technique: Measures heat changes in chemical reactions.
• Devices: Bomb calorimeters for constant volume processes.

Spontaneous and Non-Spontaneous Processes

• Direction: Natural processes tend to increase entropy.
• First Law Limitation: Cannot predict spontaneity.

Entropy and Second Law of Thermodynamics

• Entropy ($S$): Measure of disorder/randomness.
• Second Law: Total entropy change determines spontaneity.
• Equation: $\Delta S = \frac{q_{\text{rev}}}{T}$

Gibbs Energy, Spontaneity, and Equilibrium Constant

• Gibbs Energy ($G$): $G = H - TS$
• Criteria:
  • $\Delta G < 0$: Spontaneous
  • $\Delta G = 0$: Equilibrium
  • $\Delta G > 0$: Non-spontaneous
• Equilibrium Constant: $\Delta G^\circ = -RT \ln K$

Tables

• Standard Molar Enthalpies of Formation:

  Substance	$\Delta_f H^\circ$ (kJ/mol)
  $H_2O(l)$	-285.8
  $CH_4(g)$	-74.81
  $C_2H_5OH(l)$	-277.7

• Mean Bond Enthalpies:

  Bond	Enthalpy (kJ/mol)
  $H-H$	435
  $C-H$	416
  $O=O$	498

• Lattice Enthalpy Example:

  Reaction	$\Delta H^\circ$ (kJ/mol)
  $NaCl(s) \rightarrow Na^+(g) + Cl^-(g)$	+788