Advertisements
Advertisements
Question
What is the equilibrium constant Keq for the following reaction at 400 K?
\[\ce{2NOCl_{(g)} ⇌ 2NO_{(g)} + Cl2_{(g)}}\], given that H0 = 77.2 kJ mol−1 and ∆S0 = 122 JK−1 mol−1
Solution
T = 400 K; ∆H0 = 77.2 kJ mol−1 = 77200 J mol−1;
∆S0 = 122 JK−1 mol−1
∆G0 = −2.303 RT log Keq
log Keq = `(-∆"G"^0)/(2.303 "RT")`
log Keq = `((-∆"H"^0 - "T"∆"S"^0))/(2.303 "RT")`
log Keq = `-((77200 - 400 xx 122)/(2.303 xx 8.314 xx 400))`
log Keq = `-((28400)/7659)`
log Keq = − 3.7080
Keq = antilog (−3.7080)
Keq = 1.95 × 10−4
APPEARS IN
RELATED QUESTIONS
Define is Gibb’s free energy.
The equilibrium constant of a reaction is 10, what will be the sign of ∆G? Will this reaction be spontaneous?
What are the conditions for the spontaneity of a process?
List the characteristics of Gibbs free energy.
For the reaction \[\ce{Ag2O_{(s)} -> 2Ag_{(s)} + 1/2O2_{(g)}}\]: ΔH = 30.56 kJ mol−1 and ΔS = 6.66 JK−1 mol−1 (at 1 atm). Calculate the temperature at which ΔG is equal to zero. Also predict the direction of the reaction
- at this temperature and
- below this temperature.
When 1-pentyne (A) is treated with 4N alcoholic KOH at 175°C, it is converted slowly into an equilibrium mixture of 1.3% 1-pentyne (A), 95.2% 2-pentyne (B) and 3.5% of 1, 2 pentadiene (C) the equilibrium was maintained at 175°C, calculate ΔG0 for the following equilibria.
\[\ce{B ⇌ A}\] `Δ"G"_1^0` = ?
\[\ce{B ⇌ C}\] `Δ"G"_2^0` = ?
At 33K, N2O4 is fifty percent dissociated. Calculate the standard free energy change at this temperature and at one atmosphere.
Find out the value of the equilibrium constant for the following reaction at 298 K, \[\ce{2NH3_{(g)} + CO2_{(g)} ⇌ NH2CONH2_{(aq)} + H2O_{(l)}}\] Standard Gibbs energy change, `∆"G"_"r"^0` at the given temperature is –13.6 kJ mol−1.
