For a given reaction delta h=35.5 kJ mol-¹ and delta s= 83.6. the reac...
Spontaneous when delta G is negative
( delta G = delta H -T delta s)
so H- T S < />
H
H/S < />
~425
For a given reaction delta h=35.5 kJ mol-¹ and delta s= 83.6. the reac...
Explanation:
To determine the temperature at which a reaction becomes spontaneous, we can use the Gibbs free energy 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.
In order for a reaction to be spontaneous, ΔG must be negative. This means that the change in enthalpy and the change in entropy must be such that the term -TΔS is greater than ΔH.
In this case, we are given that ΔH = 35.5 kJ mol⁻¹ and ΔS = 83.6. Since ΔS is positive, we know that the reaction is more favorable at higher temperatures, as the term -TΔS will become more negative.
To find the temperature at which the reaction becomes spontaneous, we can set ΔG equal to zero and solve for T:
0 = ΔH - TΔS
TΔS = ΔH
T = ΔH/ΔS
Plugging in the given values:
T = 35.5 kJ mol⁻¹ / 83.6
T ≈ 0.425 K
Therefore, the reaction is spontaneous at temperatures greater than approximately 425 K.
Conclusion:
In conclusion, the reaction is not spontaneous at all temperatures because the term -TΔS becomes less negative as the temperature decreases. The reaction becomes spontaneous when the term -TΔS is greater than ΔH, and this occurs at temperatures greater than approximately 425 K.
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