For a spontaneous process, the Gibbs free energy change is negative.

The relationship is given by H = U + PV, where H is enthalpy, U is internal energy, and PV is the pressure-volume work.

The relationship is given by the equation H = U + PV, where H is enthalpy, U is internal energy, P is pressure, and V is volume.

At constant pressure, the relationship is given by ΔH = ΔU + PΔV.

A higher entropy generally indicates a spontaneous process, as spontaneous processes tend to increase the overall disorder of the system.

Entropy generally increases with increasing temperature due to increased molecular motion and disorder.

A negative ΔG (< 0) indicates that a reaction is spontaneous under the given conditions.

The relationship is given by ΔG = -RT ln(K), where R is the gas constant and T is the temperature in Kelvin.

For an ideal gas, the heat capacity at constant pressure (C_p) is greater than the heat capacity at constant volume (C_v).

For an ideal gas, C_p is always greater than C_v due to the work done during expansion.

The relationship is given by ΔG = -RT ln K, where R is the gas constant and T is temperature.

The relationship is given by ΔG = -RT ln(K), where R is the gas constant and T is the temperature in Kelvin.

The relationship is given by the equation ΔG = ΔH - TΔS, where T is the temperature in Kelvin.

At equilibrium, the change in Gibbs free energy (ΔG) is zero.

At constant pressure, the relationship is given by ΔH = ΔU + PΔV.

A negative ΔG indicates that the biochemical reaction is spontaneous under the given conditions.

Gibbs Free Energy is crucial for predicting the spontaneity and direction of a chemical reaction.

ΔGf° indicates the energy change when a compound is formed from its elements in their standard states.

The standard enthalpy change for the formation of a compound is denoted as ΔHf.

The standard enthalpy change for the formation of an element in its standard state is defined as zero.

The standard enthalpy change for the formation of water (H2 + 1/2 O2 -> H2O) is -285.83 kJ/mol.

The standard enthalpy of formation of water is -241.82 kJ/mol.

The standard enthalpy change of formation is defined as the enthalpy change when one mole of a compound is formed from its elements in their standard states.

The standard enthalpy change of formation for elements in their standard state is defined as zero.

The standard enthalpy change of formation for water at 25°C is -241.82 kJ/mol.

The reaction decreases the number of gas molecules from 3 to 2, leading to a decrease in disorder and thus a negative ΔS°.

The reaction results in a decrease in the number of gas molecules (from 3 to 2), leading to a negative change in entropy.

The reaction decreases the number of gas molecules from 3 to 2, leading to a decrease in entropy.

ΔG° depends on the specific reaction and its conditions.

Standard molar entropy is defined as the entropy of 1 mole of a substance at 25 °C and 1 atm pressure.