For a certain reaction, Delta H(not) = -76.8 kj and delta S(not) = -217 j/k. If n =3, calculate E for the reaction at 25degrees celcius.
To calculate the standard Gibbs free energy change (ΔG°) for a reaction using the equation
ΔG° = ΔH° - TΔS°
where ΔH° is the standard enthalpy change, ΔS° is the standard entropy change, and T is the temperature in Kelvin.
Before plugging in the values and calculating ΔG°, we need to convert the units of the entropy change from J/K to kJ/K because the standard enthalpy change is given in kJ.
ΔS(not) = -217 J/K = -0.217 kJ/K
Now, we can substitute the values into the equation:
ΔG° = -76.8 kJ - (25 + 273) K * (-0.217 kJ/K)
Calculating the expression inside the brackets first will give:
-(25 + 273) K * (-0.217 kJ/K) = -298 K * (-0.217 kJ/K) = 64.666 kJ
Substituting this value back into the equation:
ΔG° = -76.8 kJ - 64.666 kJ
ΔG° = -141.466 kJ
Now, we can calculate the non-standard Gibbs free energy change (ΔG) using the equation:
ΔG = ΔG° + RTln(Q)
where R is the ideal gas constant (8.314 J/(mol·K)), T is the temperature in Kelvin, and Q is the reaction quotient.
Since ΔG° = -141.466 kJ and n = 3 (according to the question), the expression becomes:
ΔG = -141.466 kJ + (8.314 J/(mol·K) * (25 + 273) K * ln(Q)
Since the reaction is at 25 degrees Celsius, we first need to convert the temperature to Kelvin:
T = 25 + 273 = 298 K
Now, substitute the values into the equation and calculate ΔG:
ΔG = -141.466 kJ + 8.314 J/(mol·K) * 298 K * ln(Q)
To calculate Q, we need additional information about the reaction, such as the concentrations of reactants and products or any equilibrium constants provided in the question. Without that information, we cannot determine the value of Q or calculate ΔG for the reaction at 25 degrees Celsius.