Thermochemistry Study Notes

Definition: A measure of the total energy of a thermodynamic system, including internal energy and the energy required to make space for it (PV work).
Symbol: Represented by H.
Change in Enthalpy (ΔH):
- ΔH = H(products) - H(reactants)
- Positive ΔH indicates an endothermic reaction (absorbs heat).
- Negative ΔH indicates an exothermic reaction (releases heat).

Zeroth Law: If two systems are in thermal equilibrium with a third system, they are in thermal equilibrium with each other.
First Law: Energy cannot be created or destroyed, only transformed.
- ΔU = Q - W (change in internal energy = heat added to the system - work done by the system).
Second Law: The total entropy of an isolated system can never decrease over time.
- Entropy tends to increase, leading to irreversibility in natural processes.
Third Law: As temperature approaches absolute zero, the entropy of a perfect crystal approaches a constant minimum.

Exothermic Reactions:
- Release heat (ΔH < 0).
- Example: Combustion of fuels.
Endothermic Reactions:
- Absorb heat (ΔH > 0).
- Example: Photosynthesis.
Standard Enthalpy of Formation (ΔHf°): Enthalpy change when 1 mole of compound is formed from its elements in their standard states.

Definition: The amount of heat required to change the temperature of a substance by one degree Celsius (or Kelvin).
Types:
- Specific Heat Capacity (s): Heat capacity per unit mass (J/kg·K).
- Molar Heat Capacity (C): Heat capacity per mole (J/mol·K).
Formula:
- Q = mcΔT (where Q = heat added, m = mass, c = specific heat, ΔT = change in temperature).

Types:
- Melting (solid to liquid): Endothermic process (ΔH > 0).
- Freezing (liquid to solid): Exothermic process (ΔH < 0).
- Vaporization (liquid to gas): Endothermic process (ΔH > 0).
- Condensation (gas to liquid): Exothermic process (ΔH < 0).
- Sublimation (solid to gas): Endothermic process (ΔH > 0).
- Deposition (gas to solid): Exothermic process (ΔH < 0).
Phase Diagrams: Graphs that show the state of a substance at various temperatures and pressures, indicating phase changes.

Enthalpy is crucial for understanding energy changes in reactions.
The laws of thermodynamics govern the principles of energy conservation and transformation.
Heat capacity and phase transitions highlight the relationship between energy and temperature changes in substances.

Enthalpy (H) quantifies total energy in a thermodynamic system, combining internal energy and pressure-volume work.
Change in Enthalpy (ΔH) is calculated as ΔH = H(products) - H(reactants).
Endothermic reactions occur with positive ΔH, indicating heat absorption; exothermic reactions have negative ΔH, indicating heat release.

Zeroth Law: Establishes that if two systems are in equilibrium with a third, they are also in equilibrium with each other.
First Law: Dictates that energy conservation occurs where ΔU = Q - W, meaning the change in internal energy equals heat added minus work done.
Second Law: States entropy in an isolated system cannot decrease, promoting irreversibility in natural processes.
Third Law: Highlights that as temperature nears absolute zero, a perfect crystal's entropy approaches a constant minimum.

Exothermic reactions release heat, exemplified by fuel combustion, with ΔH < 0.
Endothermic reactions absorb heat, such as photosynthesis, with ΔH > 0.
Standard Enthalpy of Formation (ΔHf°) refers to the enthalpy change for forming one mole of a compound from its elements in standard states.

Heat capacity measures the heat needed to raise a substance's temperature by one degree Celsius or Kelvin.
Specific Heat Capacity (s) is the heat capacity per unit mass, while Molar Heat Capacity (C) is per mole.
The heat added (Q) can be determined by the formula Q = mcΔT, encompassing mass (m), specific heat (c), and temperature change (ΔT).

Various transitions include melting (solid to liquid) and vaporization (liquid to gas), both endothermic processes with ΔH > 0.
Freezing (liquid to solid) and condensation (gas to liquid) are exothermic, indicated by ΔH < 0.
Sublimation (solid to gas) and deposition (gas to solid) also follow endothermic and exothermic principles, respectively.
Phase diagrams visually represent the states of a substance across different temperatures and pressures, marking phase change points.

Understanding enthalpy is vital for analyzing energy changes in chemical reactions.
The thermodynamic laws underpin the principles of energy conservation and transformation throughout natural processes.
The interplay between heat capacity and phase transitions emphasizes the interaction between energy and temperature in physical substances.