General Chemistry: Thermochemistry and Related Concepts Study Notes

Thermochemistry

System, Surroundings, and Types of Systems

Thermochemistry studies the energy changes, especially heat, that occur during chemical reactions and changes of state.

• System: The part of the universe under study (e.g., the reactants and products in a reaction vessel).

• Surroundings: Everything outside the system.

• Types of Systems:

  • Open System: Can exchange both matter and energy with surroundings.

  • Closed System: Can exchange energy but not matter.

  • Isolated System: Cannot exchange matter or energy.

Energy, Work, and Heat

Energy is the capacity to do work or produce heat. In thermochemistry, we focus on heat (q) and work (w).

• Work (w): Energy used to move an object against a force.

• Heat (q): Energy transferred due to temperature difference.

• Internal Energy (U): The total energy contained within a system.

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transferred or transformed.

Equation:

• $\Delta U = q + w$

Enthalpy (H)

Enthalpy is a state function that represents the heat content of a system at constant pressure.

• Definition: $H = U + PV$

• Change in Enthalpy: $\Delta H = \Delta U + P\Delta V$

• At constant pressure, $\Delta H$ equals the heat exchanged: $q_p = \Delta H$

Endothermic and Exothermic Processes

Chemical reactions can absorb or release heat.

• Endothermic: Absorbs heat ($\Delta H > 0$)

• Exothermic: Releases heat ($\Delta H < 0$)

• Example: Combustion of methane is exothermic.

Calorimetry

Calorimetry is the measurement of heat flow in a chemical reaction.

• Calorimeter: Device used to measure heat changes.

• Specific Heat (c): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Equation: $q = mc\Delta T$

• Example: Calculating heat absorbed by water when heated.

Bomb Calorimetry

Used to measure heat of combustion at constant volume.

• Equation: $q_{v} = C_{cal} \Delta T$

• Where: $C_{cal}$ is the calorimeter's heat capacity.

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in.

• Equation: $\Delta H_{rxn} = \sum \Delta H_{steps}$

• Application: Used to calculate enthalpy changes for reactions not easily measured directly.

Standard Enthalpy of Formation

The standard enthalpy of formation ($\Delta H_f^\circ$) is the enthalpy change when one mole of a compound is formed from its elements in their standard states.

• Equation: $\Delta H_{rxn}^\circ = \sum n \Delta H_f^\circ (\text{products}) - \sum n \Delta H_f^\circ (\text{reactants})$

• Example: Formation of water from hydrogen and oxygen.

Bond Enthalpy

Bond enthalpy is the energy required to break one mole of a particular type of bond in a gaseous molecule.

• Equation: $\Delta H_{rxn} = \sum \text{Bond Energies (bonds broken)} - \sum \text{Bond Energies (bonds formed)}$

• Application: Estimating enthalpy changes for reactions using average bond energies.

State Functions vs. Path Functions

State functions depend only on the initial and final states, not on the path taken. Path functions depend on the process.

• State Functions: Enthalpy (H), Internal Energy (U), Pressure (P), Volume (V), Temperature (T)

• Path Functions: Work (w), Heat (q)

Tables: Comparison of System Types

The following table summarizes the differences between open, closed, and isolated systems.

System Type	Matter Exchange	Energy Exchange
Open	Yes	Yes
Closed	No	Yes
Isolated	No	No

Tables: Comparison of Endothermic and Exothermic Processes

Process Type	Heat Flow	Sign of $\Delta H$	Example
Endothermic	Absorbed	Positive	Melting ice
Exothermic	Released	Negative	Combustion

Additional info:

• Some equations and definitions have been expanded for clarity.

• Examples have been added to illustrate key concepts.