Key Points

In thermodynamics, the system is the part of the universe under observation, while the surroundings are everything else. The system and surroundings together constitute the universe.

Systems are classified based on energy and matter exchange. An open system exchanges both energy and matter, a closed system exchanges only energy, and an isolated system exchanges neither.

The first law is the law of conservation of energy, stated mathematically as $\Delta U = q + w$. Here, $\Delta U$ is the change in internal energy, q is heat absorbed by the system, and w is work done on the system.

The work done during the expansion or compression of a gas against a constant external pressure is given by $w = -P_{ex} \Delta V$. For a reversible isothermal process, the work done is $w_{rev} = -2.303 nRT \log \frac{V_f}{V_i}$.

Enthalpy is a state function defined as $H = U + pV$. The change in enthalpy ($\Delta H$) is equal to the heat absorbed or evolved at constant pressure, $\Delta H = q_p$.

The relationship between the change in enthalpy and the change in internal energy for reactions involving gases is given by the equation $\Delta H = \Delta U + \Delta n_g RT$, where $\Delta n_g$ is the change in the number of moles of gaseous products and reactants.

Extensive properties depend on the amount of matter (e.g., mass, volume, enthalpy), while intensive properties do not (e.g., temperature, pressure, density).

The standard enthalpy of formation ($\Delta_f H^\ominus$) is the enthalpy change when one mole of a substance is formed from its constituent elements in their most stable reference states. By convention, $\Delta_f H^\ominus$ of an element in its reference state is zero.

Hess's Law states that the total enthalpy change for a reaction is the same, whether the reaction takes place in one step or in a series of steps. It allows for the calculation of reaction enthalpies that cannot be measured directly.

The standard enthalpy of a reaction ($\Delta_r H^\ominus$) can be calculated from standard enthalpies of formation: $\Delta_r H^\ominus = \sum \Delta_f H^\ominus (\text{products}) - \sum \Delta_f H^\ominus (\text{reactants})$.

Bond enthalpy is the energy required to break one mole of a specific type of bond in gaseous molecules. Reaction enthalpy can be estimated using bond enthalpies: $\Delta_r H^\ominus = \sum (\text{bond enthalpies})_{reactants} - \sum (\text{bond enthalpies})_{products}$.

A spontaneous process is one that has the potential to proceed on its own without external assistance. Spontaneity is determined by the total entropy change of the system and surroundings, not just by a decrease in enthalpy.

Entropy is a thermodynamic state function that measures the degree of randomness or disorder of a system. For any spontaneous process, the total entropy of the universe increases: $\Delta S_{total} = \Delta S_{system} + \Delta S_{surroundings} > 0$.

Gibbs free energy combines enthalpy and entropy into a single value to predict spontaneity. The Gibbs equation is $\Delta G = \Delta H - T\Delta S$.

For a process at constant temperature and pressure: if $\Delta G < 0$, the process is spontaneous; if $\Delta G > 0$, the process is non-spontaneous; if $\Delta G = 0$, the system is at equilibrium.

The standard Gibbs free energy change of a reaction is related to its equilibrium constant (K) by the equation $\Delta_r G^\ominus = -RT \ln K$. This can also be written as $\Delta_r G^\ominus = -2.303 RT \log K$.