So I have had several homework problem like this and managed to solve them easily, but these two problems I cannot get the correct answer for.
Using a Table of thermodynamic data, calculate the change in Gibbs free energy for each of the following reactions. In each case indicate whether the reaction is spontaneous under standard conditions. Here are some more delta G values: N2H4 (g) = 159.4 kJ/mol, FeO (s) = -255.2 kJ/mol, FeCl3 (s) = -334 kJ/mol, H3PO4 (aq) = -1142.6 kJ/mol
2 NH3(g) N2H4(g) + H2(g)
i got to be 187kj but its wrong.
2 CH3OH(l) + 3 O2(g) 2 CO2(g) + 4 H2O(l)
I got -1693 and its wrong.
I figured it out. I used the deltaH values in the chart instead of the DeltaG value.
To calculate the change in Gibbs free energy (ΔG) for a reaction, you need to use the thermodynamic data, particularly the ΔG values for each component involved in the reaction. From the given values, we have:
1. For the reaction: 2 NH3(g) → N2H4(g) + H2(g)
The ΔG value for N2H4(g) is 159.4 kJ/mol. The ΔG value for H2(g) is not provided, so we need to calculate it using the ΔG values of other reactions involving H2(g). One such reaction is the formation of H2(g) from its constituent elements:
H2(g) → 2 H(g)
The ΔG value for this reaction is 0 kJ/mol. As there are two moles of H2(g) in the desired reaction, the ΔG value for H2(g) in the desired reaction would also be 0 kJ/mol.
Adding the ΔG values for N2H4(g) and H2(g):
ΔG = ΔG(N2H4) + ΔG(H2)
= 159.4 kJ/mol + 0 kJ/mol
= 159.4 kJ/mol
Therefore, the change in Gibbs free energy for this reaction is 159.4 kJ/mol.
Now, to determine whether the reaction is spontaneous under standard conditions, we need to check the overall sign of the ΔG value. If ΔG is negative, the reaction is spontaneous; if ΔG is positive, the reaction is non-spontaneous.
In this case, ΔG is positive (159.4 kJ/mol), meaning the reaction is non-spontaneous under standard conditions.
2. For the reaction: 2 CH3OH(l) + 3 O2(g) → 2 CO2(g) + 4 H2O(l)
We need to calculate the ΔG value for CO2(g) using the given ΔG values. The ΔG value for CO2(g) is not provided, so again we need to use other reactions involving CO2(g) to calculate it. One such reaction is the formation of CO2(g) from its constituent elements:
C(s) + O2(g) → CO2(g)
The ΔG value for this reaction is -394 kJ/mol. As there are two moles of CO2(g) in the desired reaction, the ΔG value for CO2(g) in the desired reaction would be -2 * 394 kJ/mol = -788 kJ/mol.
Now, adding the ΔG values for CO2(g) and H2O(l):
ΔG = ΔG(CO2) + ΔG(H2O)
= -788 kJ/mol + 0 kJ/mol
= -788 kJ/mol
Therefore, the change in Gibbs free energy for this reaction is -788 kJ/mol.
Since the ΔG value is negative (-788 kJ/mol), the reaction is spontaneous under standard conditions.
It's important to note that the values provided here are specific to the problem given. Always ensure that the thermodynamic data you are using matches the conditions and substances involved in the problem.