The following reaction has the following standard thermodynamic parameters:

Reaction: C(g) + D(g)----> E(g)
Delta H rxn=-26.1kJ/mol and
Delta S rxn=-69.8J/(mol*K).

Calculate the temperature at which the reaction becomes nonspontaneous.

See the earlier post.

To determine the temperature at which the reaction becomes nonspontaneous, we need to use the relationship between Gibbs free energy change (ΔG), enthalpy change (ΔH), and entropy change (ΔS) given by the equation:

ΔG = ΔH - TΔS

In this equation, T represents the temperature in Kelvin. At the temperature where the reaction becomes nonspontaneous, ΔG will be equal to zero.

Since we are given the values of ΔH and ΔS, we can substitute them into the equation:

0 = (-26.1 kJ/mol) - T * (-69.8 J/(mol*K))

Notice that we need to convert the units so that they are consistent. Since ΔH is given in kJ/mol and ΔS is given in J/(mol*K), we need to convert ΔH to J/mol:

ΔH = -26.1 kJ/mol * (1000 J/1 kJ) = -26,100 J/mol

Now we can rearrange the equation and solve for T:

TΔS = -26,100 J/mol
ΔS = -69.8 J/(mol*K)

T = -26,100 J/mol / -69.8 J/(mol*K)

T = 374 K

Therefore, the reaction becomes nonspontaneous at a temperature of 374 Kelvin.