To determine whether the reaction is spontaneous, we can use the Gibbs free energy equation, ( \Delta G = \Delta H - T\Delta S ). For the reaction to be spontaneous, ( \Delta G ) must be less than 0. Given ( \Delta H = -92 , \text{kJ/mol} ) and ( \Delta S = -0.199 , \text{kJ/(mol K)} ), we can set up the inequality ( -92 , \text{kJ/mol} - T(-0.199 , \text{kJ/(mol K)}) < 0 ). Solving this will give the temperature threshold above which the reaction becomes spontaneous.
It is not spontaneous.
3600 K
To determine the temperature at which the reaction becomes spontaneous, we can use the Gibbs free energy equation: ΔG = ΔH - TΔS. A reaction is spontaneous when ΔG < 0. Given ΔH = -92 kJ/mol and ΔS = -0.199 kJ/(mol·K), we set ΔG to 0 and solve for T: 0 = -92 kJ/mol - T(-0.199 kJ/(mol·K)). This simplifies to T = 462.31 K. Thus, the reaction is spontaneous at temperatures above approximately 462 K.
it can never be spontanious
The reaction 2NH3(g) ⇌ N2(g) + 3H2(g) has a ΔH of +92.4 kJ/mol, indicating that it is endothermic, meaning it absorbs heat from the surroundings. This suggests that increasing the temperature would favor the formation of products (N2 and H2) according to Le Chatelier's principle. Additionally, the reaction involves a decrease in the number of moles of gas, which may affect the equilibrium position depending on the pressure applied.
The condition for a reaction to be spontaneous is ΔG < 0, where ΔG = ΔH - TΔS. At the temperature where ΔG becomes negative, the reaction will be spontaneous. You can calculate this temperature using the given values of ΔH and ΔS.
I2(s) --> I2(g); dH=62.4kJ/mol; dS=0.145kJ/mol. The reaction will favor the product at this temperature. Your entropy is positive and your enthalpy is also positive, so this reaction will not be spontaneous at all temperatures. But at room temperature, which is 298K, it will be spontaneous and proceed left to right. (this is the sublimation of iodine)
it is never spontaneous
It is not spontaneous.
3600 K
To determine the temperature at which the reaction becomes spontaneous, we can use the Gibbs free energy equation: ΔG = ΔH - TΔS. A reaction is spontaneous when ΔG < 0. Given ΔH = -92 kJ/mol and ΔS = -0.199 kJ/(mol·K), we set ΔG to 0 and solve for T: 0 = -92 kJ/mol - T(-0.199 kJ/(mol·K)). This simplifies to T = 462.31 K. Thus, the reaction is spontaneous at temperatures above approximately 462 K.
it can never be spontanious
G= 0 kJ/mol
The reaction is exothermic because the enthalpy change is negative (-890 kJ/mol). The reaction may be spontaneous at low temperatures due to the negative entropy change (-0.24 kJ/(mol K)), which decreases the overall spontaneity of the reaction.
it can never be spontanious
At 500 K, the reaction will favour the formation of gaseous I2 since the positive change in enthalpy indicates the reaction is endothermic. The positive change in entropy suggests an increase in disorder, further favoring the formation of gaseous I2 at higher temperatures.
The reaction 2NH3(g) ⇌ N2(g) + 3H2(g) has a ΔH of +92.4 kJ/mol, indicating that it is endothermic, meaning it absorbs heat from the surroundings. This suggests that increasing the temperature would favor the formation of products (N2 and H2) according to Le Chatelier's principle. Additionally, the reaction involves a decrease in the number of moles of gas, which may affect the equilibrium position depending on the pressure applied.