Question

For the reaction A + B → C + D, ΔH° = +40 kJ and ΔS° = +50 J/K. Therefore, the reaction under standard conditions is: A) spontaneous only at temperatures between 10 K and 800 K. B) spontaneous at temperatures greater than 800 K. C) nonspontaneous at all temperatures. D) spontaneous at all temperatures. E) spontaneous at temperatures less than 10 K.

Answer 1

For the reaction A + B → C + D, with ΔH° = +40 kJ and ΔS° = +50 J/K, the reaction is spontaneous at temperatures greater than 800 K as the positive ΔH° and ΔS° indicates that an increase in temperature will lead to a negative Gibbs free energy change, indicating spontaneity.

Step-by-step explanation:

The question is determining the spontaneity of the reaction A + B → C + D with a given enthalpy change (ΔH°) and entropy change (ΔS°). Spontaneity is assessed by the Gibbs free energy change (ΔG), which is calculated by the equation ΔG = ΔH° - TΔS°, where T is the temperature in Kelvin.

Given that ΔH° = +40 kJ and ΔS° = +50 J/K, we need to determine the temperature at which ΔG becomes negative, as this indicates spontaneity. Since ΔH° is positive and ΔS° is also positive, the spontaneity depends on the temperature. Here we can use the formula:

ΔG = ΔH° - TΔS°

Spontaneity requires that ΔG < 0. Therefore, we solve for T:

0 > ΔH° - TΔS°
ΔH° < TΔS°
T > ΔH°/ΔS°
T > 40,000 J / 50 J/K
T > 800 K

Thus, the reaction in question is spontaneous at temperatures greater than 800 K, because it is only above this temperature that ΔG turns negative, making choice (B) the correct answer.

Answer 2

spontaneous at temperatures greater than 800 K.

Step-by-step explanation:

From the thermodynamic equation ,

ΔG = ΔH - TΔS

For spontaneous reaction ΔG = - ve

ΔH - TΔS < 0

TΔS > ΔH

T > ΔH / ΔS

> 40000 / 50

> 800 K

At temperature greater than 800K , reaction will be spontaneous.