General Chemistry
The Gibbs free energy is the energy available to do work in a system (specifically, a closed system at constant temperature and pressure). The change in Gibbs free energy, ΔG, informs us whether free energy is absorbed or released during a reaction. ΔG can be calculated by subtracting the product of temperature and change in entropy (T * ΔS) from the change in enthalpy (ΔH). When ΔG is negative, a reaction is exergonic, which means it occurs spontaneously and releases energy. On the other hand, if ΔG is positive, a reaction is endergonic, non-spontaneous, and absorbs energy.
An important concept related to Gibbs free energy is the standard Gibbs free energy of formation, which is the ΔG that occurs when one mole of a substance is formed at 298 Kelvin and one atmosphere of pressure, and the reacting elements are in their standard states of matter.
Lesson Outline
<ul> <li>Introduction to Gibbs free energy and its role in reactions</li> <ul> <li>Energy available to do work in a system</li> <li>Change in Gibbs free energy (ΔG) indicates energy absorbed or released during a reaction</li> </ul> <li>Calculating ΔG</li> <ul> <li>Formula: ΔG = ΔH - T * ΔS</li> <li>ΔH: change in enthalpy or heat inside a system during a reaction</li> <li>T: temperature of reaction in Kelvin</li> <li>ΔS: change in molecular disorder (entropy) during reaction</li> </ul> <li>Exergonic and endergonic reactions</li> <ul> <li>Negative ΔG indicates an exergonic reaction</li> <ul> <li>Energy is released</li> <li>Reaction happens spontaneously</li> <li>Spontaneity does not always mean a fast reaction</li> </ul> <li>Positive ΔG indicates an endergonic reaction</li> <ul> <li>Energy is absorbed, such as heat</li> <li>Reaction is non-spontaneous and requires external energy</li> </ul> </ul> <li>Standard Gibbs free energy of formation</li> <ul> <li>ΔG when one mole of a substance is formed</li> <li>Conditions: 298 Kelvin, one atmosphere of pressure</li> <li>Reacting elements must be in their standard states of matter</li> </ul> </ul>
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FAQs
The relationship between standard Gibbs free energy (ΔG°) and equilibrium is expressed through the equilibrium constant (k) via the equation ΔG° = -RT ln k, where R is the universal gas constant, and T is the temperature in Kelvin. If ΔG° is negative, the reaction is spontaneous and will proceed towards equilibrium, with k > 1. If ΔG° is positive, the reaction is non-spontaneous and will not proceed towards equilibrium unless the conditions change, with k < 1.
Exergonic and endergonic processes are characterized by their effect on the standard Gibbs free energy. Exergonic processes have a negative ΔG, indicating that energy is released during the process, making it spontaneous. In contrast, endergonic processes have a positive ΔG, signifying that energy is absorbed, and so they are non-spontaneous. ΔG values can be calculated using the equation ΔG = ΔH - TΔS, where ΔH is the change in enthalpy, ΔS is the change in entropy, and T is the temperature in Kelvin.
The spontaneity of a reaction is determined by its Gibbs free energy change (ΔG). According to the equation ΔG = ΔH - TΔS, both the change in enthalpy (ΔH) and change in entropy (ΔS) play crucial roles. If ΔH is negative (exothermic) and ΔS is positive, the reaction will always be spontaneous. Conversely, if ΔH is positive (endothermic) and ΔS is negative, the reaction will be non-spontaneous. However, if ΔH and ΔS both have either positive or negative values, the temperature will dictate the spontaneity of the reaction.
Standard Gibbs free energy of formation (ΔGf°) values are tabulated quantities representing the change in Gibbs free energy when one mole of a substance is formed under standard conditions from its constituent elements. They help predict the spontaneity of a reaction by allowing one to calculate the overall Gibbs free energy change (ΔG) using the equation ΔG = ΣΔGf°(products) - ΣΔGf°(reactants). A negative ΔG indicates a spontaneous reaction, while a positive ΔG indicates a non-spontaneous reaction.
