CH1401 Thermodynamics Lecture 3

Overview of Thermodynamics Concepts

1. Energy and Work

• Energy can be identified in forms such as work (W) and heat (q), influencing thermodynamic processes.

• Final and initial energy values are significant in determining the change in energy in a system.

2. Heat (q)

2.1 Definition

• Heat refers to the energy transferred to or from a system due to a temperature difference.

2.2 Temperature-Heat Relationship

• For an ideal gas, temperature increase is proportional to the heat supplied.

• Heat capacity (C) characterizes this relationship as: C = q/ delta T

  • Units: J K⁻¹

  • Values vary for different phases of water: ice, liquid, steam.

2.3 Latent Heat

• Temperature change may not occur during phase changes, where latent heat is absorbed.

3. Heat Capacity

3.1 Types

• Extensive Property: Depends on the amount of substance.

• Intensive Property: Independent of the amount of substance.

  • Specific Heat Capacity (Cs):[ C_s = \frac{q}{m \Delta T} ]

  • Units: J K⁻¹ g⁻¹

  • Molar Heat Capacity (Cm):[ C_m = \frac{q}{n \Delta T} ]

  • Units: J K⁻¹ mol⁻¹

3.2 Worked Example

• Calculate energy to raise the temperature of 2 moles of liquid benzene by 20 K using tabulated values for Cs and Cm.

4. Internal Energy (U)

4.1 Definition

• U accounts for all kinetic and potential energy in a system, comprising contributions from atomic, ionic, and molecular motions.

4.2 Change in Internal Energy

• The formula is given by:[ \Delta U = q + w ]

  • Where: q = heat added, w = work done on/by the system.

4.3 Path Dependency

• Change in internal energy is a state function, implying it only depends on initial and final states, not the path taken.

5. First Law of Thermodynamics

• States that internal energy in an isolated system remains constant; reinforces energy conservation.

• Energy changes (q and w) should sum to zero for a closed system.

6. Enthalpy (H)

6.1 Definition

• Enthalpy accounts for heat changes at constant pressure:[ H = U + pV ]

6.2 Enthalpy Change

• Given by:[ \Delta H = \Delta U + p \Delta V ]

• Heat transfer under constant pressure is equal to change in enthalpy.

6.3 Exothermic vs Endothermic Reactions

• Exothermic (q < 0) = Release of heat, decrease in enthalpy.

• Endothermic (q > 0) = Absorption of heat, increase in enthalpy.

7. Calculations and Examples

• Calculating heat capacity from calorimetry data.

• Application of heat transfer equations in phase transitions and state changes.

8. Self-test Questions

• Focus on calculations involving specific heat capacity and internal energy change.

• Example questions provide practical engagement with thermodynamic principles.

9. Review Points

• Understand concepts of heat capacity, specific and molar heat capacities.

• Internal energy as a state function and its implications in thermodynamic processes.

• Grasp the first law of thermodynamics and enthalpy definitions and applications.