Thermodynamics: Laws, Enthalpy, Entropy, and Gibbs Free Energy
Thermodynamic Processes
A thermodynamic process is the transformation of a system from an initial state of equilibrium to a final state of equilibrium. A system is in equilibrium when it meets the following conditions:
- Chemical Equilibrium: The composition does not change.
- Mechanical Equilibrium: No macroscopically observable movements occur.
- Thermal Equilibrium: The temperature is uniform throughout the system.
These changes can be:
- Reversible: The transformation occurs through a succession of equilibrium states, allowing a return to any previous state.
- Irreversible: The transformation is not always in equilibrium, and the system cannot return to a previous state.
Types of transformations include:
- Adiabatic: Q = 0
- Isothermal: Temperature is constant.
- Isometric: Volume is constant.
- Isobaric: Pressure is constant.
First Law of Thermodynamics
The first law states that energy cannot be created or destroyed, only transformed. In thermodynamics, energy is classified as:
- Heat (Q): Energy transferred due to a temperature difference.
- Work (W): Any energy transfer other than heat.
The first law is expressed as: ΔU = Q + W, where ΔU is the change in internal energy. Q and W are positive when energy is added to the system.
Enthalpy
Enthalpy (H) is defined as H = Q – VΔP. It accounts for energy losses as heat (Q) in processes where volume is not constant. Enthalpy helps determine if a reaction is exothermic (H < 0) or endothermic (H > 0).
Applications of the First Law
- Isothermal Process: Temperature is constant, so Q = W.
- Isometric Process: Volume is constant, so ΔH = Q.
- Isobaric Process: Pressure is constant, so ΔH = Qp.
- Adiabatic Process: Heat is zero, so ΔU = W.
Chemical and Thermochemical Reactions
Chemical reactions involve breaking and forming bonds, requiring energy. If the energy required to break bonds is less than the energy released when forming new bonds, the reaction is exothermic. Otherwise, it is endothermic.
A thermochemical equation specifies the physical states of substances and the conditions (pressure and temperature) under which the reaction occurs. Standard conditions are often used, with a standard enthalpy of formation defined for elements in their most stable form. Key standard enthalpies include enthalpy of formation and enthalpy of combustion.
Bond Enthalpy
Forming a product involves breaking existing bonds and forming new ones. This requires energy consumption for bond breaking and energy release for bond formation.
Hess’s Law
Enthalpy is a state function, meaning it depends only on the initial and final states. If a reaction occurs in multiple stages, the total enthalpy change is the sum of the enthalpy changes of the intermediate reactions.
Second Law of Thermodynamics: Entropy
A spontaneous process occurs without external interaction. These processes reach thermodynamic equilibrium and are irreversible. They increase disorder, measured by a state function called entropy (S). Entropy change (ΔS) is the difference between the entropy of the products and the entropy of the reactants, multiplied by the number of moles.
As entropy is a state function, we can only measure the change in entropy.
Gibbs Free Energy
Gibbs free energy (G) relates enthalpy and entropy change. The Gibbs free energy of formation (ΔG°f) is the change in free energy when synthesizing 1 mole of a compound from its elements in their standard states.
ΔG determines reaction spontaneity. Spontaneous (exothermic) reactions have ΔG < 0, while non-spontaneous (endothermic) reactions have ΔG > 0. ΔG is the primary criterion for spontaneity.
Gibbs Free Energy (repeated section)
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