$ΔG^°$ is the standard Gibbs free energy change, which is a fundamental concept in chemistry that describes the spontaneity and feasibility of a chemical reaction. It combines the effects of enthalpy (energy released or absorbed) and entropy (disorder or randomness) to determine the overall energy change and the direction a reaction will naturally proceed. $ΔG^°$ is crucial in understanding potential, free energy, and equilibrium in chemical systems.
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$ΔG^°$ is calculated as $ΔG^° = ΔH^° - TΔS^°$, where $ΔH^°$ is the standard enthalpy change and $ΔS^°$ is the standard entropy change.
If $ΔG^° < 0$, the reaction is spontaneous and releases free energy. If $ΔG^° > 0$, the reaction is non-spontaneous and requires an input of free energy.
The magnitude of $ΔG^°$ indicates the driving force of the reaction, with more negative values corresponding to a greater tendency for the reaction to occur.
At equilibrium, $ΔG^° = 0$, meaning the forward and reverse reaction rates are equal and there is no net change in the system.
$ΔG^°$ is related to the equilibrium constant ($K_c$) through the equation $ΔG^° = -RT\ln K_c$, where R is the gas constant and T is the absolute temperature.
Review Questions
Explain how $ΔG^°$ is calculated and how its value relates to the spontaneity of a reaction.
$ΔG^°$ is calculated as $ΔG^° = ΔH^° - TΔS^°$, where $ΔH^°$ is the standard enthalpy change and $ΔS^°$ is the standard entropy change. If $ΔG^° < 0$, the reaction is spontaneous and releases free energy, meaning it will occur naturally. If $ΔG^° > 0$, the reaction is non-spontaneous and requires an input of free energy to proceed. The magnitude of $ΔG^°$ indicates the driving force of the reaction, with more negative values corresponding to a greater tendency for the reaction to occur.
Describe the relationship between $ΔG^°$ and the equilibrium constant ($K_c$) of a reaction.
The relationship between $ΔG^°$ and the equilibrium constant ($K_c$) is given by the equation $ΔG^° = -RT\ln K_c$, where R is the gas constant and T is the absolute temperature. This means that the value of $ΔG^°$ can be used to determine the equilibrium constant of a reaction, and vice versa. At equilibrium, when $ΔG^° = 0$, the forward and reverse reaction rates are equal, and there is no net change in the system.
Analyze how the values of $ΔH^°$ and $ΔS^°$ contribute to the overall $ΔG^°$ of a reaction and the resulting spontaneity or non-spontaneity.
The values of $ΔH^°$ and $ΔS^°$ both contribute to the overall $ΔG^°$ of a reaction through the equation $ΔG^° = ΔH^° - TΔS^°$. Reactions with a negative $ΔH^°$ (exothermic) and a positive $ΔS^°$ (increase in disorder) will have a negative $ΔG^°$, making the reaction spontaneous. Conversely, reactions with a positive $ΔH^°$ (endothermic) and a negative $ΔS^°$ (decrease in disorder) will have a positive $ΔG^°$, making the reaction non-spontaneous and requiring an input of free energy to proceed.
Related terms
Enthalpy ($ΔH^°$): Enthalpy is the measure of the total energy of a system, including the energy released or absorbed during a chemical reaction at constant pressure.
Entropy ($ΔS^°$): Entropy is a measure of the disorder or randomness of a system. Reactions that increase the entropy of the universe are more likely to occur spontaneously.
Equilibrium is the state where the forward and reverse reaction rates are equal, resulting in no net change in the concentrations of reactants and products.