Chemistry Thermodynamics Concepts

Key Concepts in Chemistry Thermodynamics

Thermodynamics: Study of energy, heat, and work in chemical systems.
Laws of Thermodynamics:
1. Zeroth Law: If two systems are each in thermal equilibrium with a third system, they are in thermal equilibrium with each other.
2. First Law: Energy cannot be created or destroyed, only transformed (ΔU = Q - W, where ΔU is the change in internal energy, Q is heat added to the system, and W is work done by the system).
3. Second Law: Entropy of an isolated system always increases; natural processes favor the direction of increasing disorder.
4. Third Law: As the temperature approaches absolute zero, the entropy of a perfect crystal approaches zero.
Key Terms:
- System: The part of the universe being studied (open, closed, isolated).
- Surroundings: Everything outside the system.
- State Functions: Properties that depend only on the state of the system (e.g., internal energy, enthalpy, entropy).
- Path Functions: Properties that depend on the path taken (e.g., heat, work).
Types of Processes:
- Isothermal: Constant temperature (ΔU = 0).
- Adiabatic: No heat exchange (Q = 0).
- Isochoric: Constant volume (W = 0).
- Isobaric: Constant pressure.
Enthalpy (H): Measure of total heat content; H = U + PV (where P is pressure and V is volume).
Entropy (S): Measure of disorder or randomness in a system.
Gibbs Free Energy (G): Indicates the spontaneity of a process; G = H - TS (where T is temperature in Kelvin).
- If ΔG < 0, the process is spontaneous.
- If ΔG > 0, the process is non-spontaneous.
Heat Capacity (C): Amount of heat required to change the temperature of a substance by one degree.
- Specific heat capacity (c): Heat capacity per unit mass (C = mc).
Calorimetry: Measurement of heat changes in chemical reactions; uses calorimeters to measure heat (Q = mcΔT).

Applications:

Understanding reaction spontaneity, equilibrium, and chemical reaction rates.
Designing energy-efficient processes in chemical engineering.
Predicting physical changes and properties of materials.

Thermodynamics

Study of energy and its transformations in chemical systems.
Four Fundamental Laws: Define the behavior of energy in the universe.
Zeroth Law: Systems in thermal equilibrium with the same third system are also in equilibrium with each other.
First Law: Energy cannot be created or destroyed, only transformed.
- Expressed through the equation: ΔU = Q - W
  - ΔU represents the change in internal energy.
  - Q represents heat added to the system.
  - W represents work done by the system.
Second Law: Entropy of an isolated system always increases over time.
- Natural processes tend toward greater disorder.
Third Law: As the temperature of a crystalline substance approaches absolute zero, its entropy approaches zero.

Key Terms

System: The specific part of the universe under observation. Can be:
- Open System: Can exchange both energy and matter with its surroundings.
- Closed System: Exchanges energy but not matter.
- Isolated System: Does not exchange energy or matter.
Surroundings: Everything outside the system.
State Functions: Properties that depend only on the current state of the system, not the path taken to reach that state.
- Examples: Internal energy (U), enthalpy (H), entropy (S).
Path Functions: Properties that depend on the specific path taken to reach a state.
- Examples: Heat (Q) and work (W).

Types of Processes

Isothermal: Constant temperature (ΔU = 0).
Adiabatic: No heat exchange (Q = 0).
Isochoric: Constant volume (W = 0).
Isobaric: Constant pressure.

Enthalpy

Measure of total heat content of a system.
H = U + PV
- P is pressure.
- V is volume.

Entropy

Measure of disorder or randomness in a system.
A high entropy state is very disordered.

Gibbs Free Energy

Indicates the spontaneity of a process.
G = H - TS
- T is temperature in Kelvin.
Spontaneous processes: ΔG < 0.
Non-spontaneous processes: ΔG > 0.

Heat Capacity

Amount of heat required to change the temperature of a substance by one degree.
C = mc (where c is specific heat capacity and m is mass).

Calorimetry

Measuring heat changes in chemical reactions.
Uses calorimeters to measure heat.
Q = mcΔT

Applications of Thermodynamics

Understanding the spontaneity of chemical reactions and the factors that affect them.
Predicting chemical equilibrium and reaction rates.
Optimizing energy efficiency in industrial chemical processes.
Determining the physical changes and chemical properties of materials.