Gibbs free energy, denoted as ΔG, is a crucial concept in thermodynamics that quantifies the energy change associated with a chemical or physical process capable of performing work. The sign of ΔG is instrumental in determining the spontaneity of a reaction. When ΔG is less than 0 (negative), the reaction is considered spontaneous, indicating that it can occur without external input. Conversely, if ΔG is greater than 0 (positive), the reaction is non-spontaneous, meaning it requires energy to proceed. In cases where ΔG equals 0, the system is at equilibrium, signifying a state where the forward and reverse reactions occur at the same rate, and there is no net change. Understanding the implications of ΔG is essential for predicting the behavior of chemical reactions and their feasibility in various conditions.
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Gibbs Free Energy (Simplified): Videos & Practice Problems
Gibbs free energy (ΔG) indicates the spontaneity of a reaction. A negative ΔG signifies a spontaneous reaction, while a positive ΔG indicates non-spontaneity. At equilibrium, ΔG equals zero. The spontaneity can also be inferred from enthalpy (ΔH) and entropy (ΔS) signs: both positive leads to spontaneity at high temperatures, while both negative does so at low temperatures. The formula for ΔG is , requiring consistent units for accurate calculations.
Gibbs Free Energy represents energy associated with a chemical reaction that can be used to do work.
Gibbs Free Energy
Gibbs Free Energy (Simplified) Concept 1
Gibbs Free Energy (Simplified) Concept 1 Video Summary
Gibbs Free Energy (Simplified) Example 1
Gibbs Free Energy (Simplified) Example 1 Video Summary
In the context of reversible reactions, the Gibbs free energy change (ΔG) provides insight into the spontaneity of a reaction. When ΔG is small and positive, it indicates that the reaction is non-spontaneous in the forward direction, meaning that the reactants are favored over the products. This small positive value suggests that the system is close to equilibrium, as ΔG approaches zero at equilibrium.
In this scenario, since the forward reaction is non-spontaneous, the reverse reaction becomes spontaneous. Therefore, while the forward reaction does not proceed readily, the reverse reaction can occur with relative ease. The proximity of ΔG to zero signifies that the system is near equilibrium, where the concentrations of reactants and products are balanced, allowing for the possibility of both forward and reverse reactions to occur.
In summary, when ΔG is small and positive, the reaction is non-spontaneous in the forward direction but spontaneous in the reverse direction, and the system is near equilibrium.
Gibbs Free Energy (Simplified) Concept 2
Gibbs Free Energy (Simplified) Concept 2 Video Summary
In thermodynamics, the spontaneity of a chemical reaction can be assessed using the signs of enthalpy change (ΔH) and entropy change (ΔS). When both ΔH and ΔS are positive, the reaction is spontaneous at high temperatures. Conversely, if both ΔH and ΔS are negative, the reaction is spontaneous at low temperatures. This indicates that temperature plays a crucial role in determining spontaneity based on the signs of these thermodynamic quantities.
In scenarios where ΔH is positive and ΔS is negative, the reaction is always non-spontaneous, regardless of temperature. On the other hand, if ΔH is negative and ΔS is positive, the reaction is spontaneous. This relationship can be summarized as follows: for reactions to be spontaneous, the combination of ΔH and ΔS must align with the temperature conditions. Understanding these principles allows for predicting the feasibility of reactions under varying thermal conditions.
Gibbs Free Energy (Simplified) Example 2
Gibbs Free Energy (Simplified) Example 2 Video Summary
In the reaction between phosphorus trichloride (PCl3) and chlorine gas (Cl2) to form phosphorus pentachloride (PCl5), the enthalpy change (ΔH) is reported as -92.50 kJ at 25 degrees Celsius. This negative value indicates that the reaction is exothermic, meaning it releases heat rather than absorbing it. Consequently, the statement claiming this is an endothermic reaction is incorrect.
When considering the equilibrium constant (K), which is defined as the ratio of products to reactants, an increase in temperature affects the position of equilibrium. Since the reaction produces one product from two reactants, the entropy change (ΔS) is negative due to the decrease in disorder. As temperature increases, the spontaneity of the reaction decreases, leading to a preference for the reactants over the products. This shift results in a decrease in the equilibrium constant (K), as the concentration of reactants increases relative to products.
To summarize the thermodynamic parameters: ΔH is negative, indicating an exothermic reaction, while ΔS is also negative, reflecting a decrease in entropy. The Gibbs free energy change (ΔG) will only be negative (indicating spontaneity) at lower temperatures, not at all temperatures. Therefore, the only accurate statement regarding this reaction is that ΔS for the reaction is negative, confirming the decrease in entropy as reactants combine to form a single product.
What are the signs of ∆H, ∆S and ∆G for the spontaneous conversion of a solid into gas?

Consider the combustion of butane gas and predict the signs of ΔS, ΔH and ∆G.
C4H10(g) + 13/2 O2(g) ⟶ 4 CO2(g) + 5 H2O(g)
Gibbs Free Energy (Simplified) Concept 3
Gibbs Free Energy (Simplified) Concept 3 Video Summary
The calculation of Gibbs free energy, denoted as ΔG, is essential in thermodynamics for understanding the spontaneity of reactions. The Gibbs free energy formula is expressed as:
\( \Delta G = \Delta H - T \Delta S \)
In this equation, ΔH represents the change in enthalpy, typically provided in kilojoules (kJ), while ΔS denotes the change in entropy, usually given in joules per Kelvin (J/K). The temperature (T) is measured in Kelvin (K). It is crucial to ensure that the units for ΔH and ΔS are consistent before performing the calculation. Since ΔH is often given in kilojoules, it is standard practice to convert ΔS from joules to kilojoules by dividing by 1000, allowing for uniformity in units.
When applying this formula, remember to include the appropriate units in your calculations to maintain accuracy. The resulting value of ΔG will indicate whether a reaction is spontaneous; a negative ΔG suggests spontaneity, while a positive ΔG indicates non-spontaneity.
Gibbs Free Energy (Simplified) Example 3
Gibbs Free Energy (Simplified) Example 3 Video Summary
In thermodynamics, the spontaneity of a reaction can be determined using the Gibbs free energy change, represented by the equation:
\( \Delta G = \Delta H - T \Delta S \)
For the given reaction, the enthalpy change (\( \Delta H \)) is -111.4 kJ, and the entropy change (\( \Delta S \)) is -25 J/K. To use these values in the equation, it is essential to convert the entropy change from joules to kilojoules. This conversion involves moving the decimal point three places to the left, resulting in:
\( \Delta S = -0.025 \, \text{kJ/K} \)
Next, substituting the values into the Gibbs free energy equation at a temperature of 298 K:
\( \Delta G = -111.4 \, \text{kJ} - (298 \, \text{K} \times -0.025 \, \text{kJ/K}) \)
Calculating the second term:
\( 298 \, \text{K} \times -0.025 \, \text{kJ/K} = -7.45 \, \text{kJ} \)
Now, substituting this back into the equation gives:
\( \Delta G = -111.4 \, \text{kJ} + 7.45 \, \text{kJ} = -103.95 \, \text{kJ} \)
Since the calculated \( \Delta G \) value is less than 0, it indicates that the reaction is spontaneous at 298 K. A negative \( \Delta G \) signifies that the reaction can proceed in the forward direction without the need for external energy input. Conversely, if \( \Delta G \) were equal to 0, the system would be at equilibrium, and if \( \Delta G \) were greater than 0, the reaction would be non-spontaneous in the forward direction, favoring the reverse reaction instead.
A particular reaction has ΔG = –350 kJ and ΔS = –350 J/K at 24°C. How much heat will be released/absorbed?
For a reaction in which ΔH = 125 kJ and ΔS = 325 J/K, determine the temperature in Celsius above which the reaction is spontaneous.
Do you want more practice?
Here’s what students ask on this topic:
What is Gibbs free energy and how does it determine the spontaneity of a reaction?
Gibbs free energy (ΔG) is a measure of the energy change in a chemical or physical process that can be used to do work. The sign of ΔG determines the spontaneity of a reaction. If ΔG is negative, the reaction is spontaneous, meaning it can proceed without external energy input. If ΔG is positive, the reaction is non-spontaneous and requires energy input to proceed. When ΔG is zero, the system is at equilibrium, indicating no net change in the reaction direction. This concept is crucial for understanding whether a reaction will occur naturally under given conditions.
How can the signs of enthalpy (ΔH) and entropy (ΔS) predict the spontaneity of a reaction?
The spontaneity of a reaction can be predicted using the signs of enthalpy (ΔH) and entropy (ΔS). If both ΔH and ΔS are positive, the reaction is spontaneous at high temperatures. If both are negative, the reaction is spontaneous at low temperatures. If ΔH is positive and ΔS is negative, the reaction is always non-spontaneous. Conversely, if ΔH is negative and ΔS is positive, the reaction is always spontaneous. These relationships help determine the conditions under which a reaction will proceed spontaneously.
What is the formula for calculating Gibbs free energy?
The formula for calculating Gibbs free energy (ΔG) is given by:
In this equation, ΔH represents the change in enthalpy, ΔS represents the change in entropy, and T is the temperature in Kelvin. It is important to ensure that the units of ΔH and ΔS are consistent, typically converting both to kilojoules (kJ) for standard calculations. This formula allows us to determine the Gibbs free energy change for a reaction, which in turn indicates its spontaneity.
Why is it important to use consistent units when calculating Gibbs free energy?
Using consistent units when calculating Gibbs free energy (ΔG) is crucial for accurate results. Typically, ΔH is given in kilojoules (kJ) and ΔS in joules per Kelvin (J/K). Since these units do not match, one must convert them to the same unit, usually kilojoules, to ensure the calculation is correct. Inconsistent units can lead to incorrect ΔG values, which would misrepresent the spontaneity of the reaction. Therefore, always convert ΔS to kJ/K by dividing by 1000 if ΔH is in kJ, ensuring the units are consistent throughout the calculation.
What does it mean when Gibbs free energy (ΔG) is zero?
When Gibbs free energy (ΔG) is zero, it means the system is at equilibrium. At this point, there is no net change in the concentrations of reactants and products, and the reaction does not favor either the forward or reverse direction. Essentially, the system is in a state of balance, and no work can be extracted from the reaction. This condition is crucial for understanding dynamic equilibrium in chemical processes, where the rates of the forward and reverse reactions are equal.