Ag(s) + Cl2(g) --> 2AgCl(s)
∆H and ∆S are 254.14kJ and –115.0J/K respectively. Calculate equilibrium constant, K for the reaction at 500K?
tbh idk
To calculate the equilibrium constant, K, for the reaction at 500K, we can use the equation:
ΔG = ΔH - TΔS
where ΔG is the change in Gibbs free energy, ΔH is the change in enthalpy, ΔS is the change in entropy, and T is the temperature in Kelvin.
First, we need to convert the value of ΔS from J/K to kJ/K by dividing by 1000:
ΔS = -115.0 J/K / 1000 = -0.115 kJ/K
Next, we can substitute the values into the equation:
ΔG = 254.14 kJ - (500 K)(-0.115 kJ/K)
= 254.14 kJ + 57.5 kJ
= 311.64 kJ
Now, we can relate ΔG to K using the equation:
ΔG = -RT ln K
where R is the gas constant (8.314 J/mol K or 0.008314 kJ/mol K) and ln represents the natural logarithm.
To get K, we need to rearrange the equation:
K = e^(-ΔG / RT)
Let's substitute the values now:
K = e^(-311.64 kJ / (0.008314 kJ/mol K)(500 K))
Calculating this on a calculator or computer program gives us:
K ≈ 2.35 x 10^73
Therefore, the equilibrium constant for the reaction at 500K is approximately 2.35 x 10^73.