Please help me with these problems, the book examples very different. I'm really confused. Any help with these will help me immensely. Thank you in advance! What I got were: 1.)c 2.)d 3.)a 4.)c 5.)b
1.) When is any endothermic reaction spontaneous?
a.)When ΔSreaction > 0
b.)When ΔSsurroundings > 0
c.)When ΔSuniverse > 0
d.)None of the above
2.) How many total microstates are possible for 2 particles that are trapped in 2 connecting flasks?
a.)1 microstate
b.)2 microstates
c.)3 microstates
d.)4 microstates
e.)5 microstates
f.)6 microstates
3.) Determine the spontaneity of this reaction:
4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(g) ΔHrxn = -1267
a.)The reaction is spontaneous at high temperatures
b.)The reaction is spontaneous at low temperatures
c.)The reaction is NOT spontaneous at any temperature.
d.)The reaction is spontaneous at all temperatures
e.)It is impossible to determine the reaction
4.) Estimate ΔG°rxn for the following reaction at 387 K.
HCN(g) + 2 H2(g) → CH3NH2(g) (ΔH°= -158.0 kJ; ΔS°= -219.9 J/K)
a.)+243 kJ
b.)-72.9 kJ
c.)+84.9 kJ
d.)-92.5 kJ
e.)-188 kJ
5.) Use the free energies of formation given below to calculate the equilibrium constant (K) for the following reaction at 298 K. [2 HNO3(aq) + NO(g) → 3 NO2(g) + H2O(l)]
(ΔG°f (kJ/mol): -110.9; 87.6; 51.3; -237.1)
a.)8.71 × 10^8
b.)0.980
c.)1.15 × 10^-9
d.)1.02
e.)5.11 × 10^-4
I think 1 is right.
I don't know about 2.
For 3 you want to know delta G for the reaction. You can calculate that from dG = dH - TdS. You have dH but you don't have T nor dS so I don't think you can calculate dG with two unknowns.
4. The answer isn't c. The mistake you made is adding dH to TdS without changing units first. Notice that dH is in kJ and TdS is in J. You must change one of them before adding; when you do that you won't get c.
5. dGo rxn = (n*dGo formation products) - (n*dGo formation reactants).
Then dG = -RT*ln*K and calculate for K.
1.) When is any endothermic reaction spontaneous?
c.) When ΔSuniverse > 0
2.) How many total microstates are possible for 2 particles that are trapped in 2 connecting flasks?
d.) 4 microstates
3.) Determine the spontaneity of this reaction:
b.) The reaction is spontaneous at low temperatures
4.) Estimate ΔG°rxn for the following reaction at 387 K.
c.) +84.9 kJ
5.) Use the free energies of formation given below to calculate the equilibrium constant (K) for the following reaction at 298 K. [2 HNO3(aq) + NO(g) → 3 NO2(g) + H2O(l)]
(ΔG°f (kJ/mol): -110.9; 87.6; 51.3; -237.1)
e.) 5.11 × 10^-4
Sure! I can help you with these problems. I'll explain each question and the steps to solve them.
1.) When is any endothermic reaction spontaneous?
To determine when an endothermic reaction is spontaneous, we need to consider the change in entropy. The correct answer is c.) When ΔSuniverse > 0. This means that the overall change in entropy of the universe is positive, which allows the reaction to occur spontaneously.
2.) How many total microstates are possible for 2 particles that are trapped in 2 connecting flasks?
To determine the total microstates, we need to calculate the total number of possible arrangements of the particles. In this case, we have two particles and two flasks. The correct answer is f.) 6 microstates. This is because each particle can be in either flask (2 options), and there are two particles, so we multiply the options together (2 x 2 = 4). However, we also need to consider that both particles could be in the same flask, so we add that possibility (4 + 1 = 5). Finally, we add the case when both flasks are empty (5 + 1 = 6).
3.) Determine the spontaneity of this reaction: 4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(g) ΔHrxn = -1267
To determine the spontaneity of the reaction, we need to consider the sign of the Gibbs free energy change (ΔG). The correct answer is a.) The reaction is spontaneous at high temperatures. The negative ΔH value indicates that the reaction is exothermic. However, we also need to consider the entropy change, which is positive because the number of gas molecules increases. At high temperatures, the positive entropy change dominates over the negative enthalpy change, making the reaction spontaneous.
4.) Estimate ΔG°rxn for the following reaction at 387 K.
HCN(g) + 2 H2(g) → CH3NH2(g) (ΔH°= -158.0 kJ; ΔS°= -219.9 J/K)
To estimate the ΔG°rxn at a given temperature, we can use the equation: ΔG°rxn = ΔH°rxn - TΔS°rxn, where T is the temperature in Kelvin. Plugging in the values, we get: ΔG°rxn = -158.0 kJ - (387 K)(-219.9 J/K). The correct answer is d.) -92.5 kJ.
5.) Use the free energies of formation given below to calculate the equilibrium constant (K) for the following reaction at 298 K. [2 HNO3(aq) + NO(g) → 3 NO2(g) + H2O(l)]
(ΔG°f (kJ/mol): -110.9; 87.6; 51.3; -237.1)
To calculate the equilibrium constant (K), we need the ΔG° values for the products and reactants. The equation is: ΔG°rxn = Σ ΔG°f(products) - Σ ΔG°f(reactants). Plugging in the values, we get: ΔG°rxn = [3(-237.1 kJ) + 51.3 kJ + 87.6 kJ] - [2(-110.9 kJ) + 2(87.6 kJ)]. Then, we can calculate K using the equation: K = e^(-ΔG°rxn / (RT)), where R is the ideal gas constant (8.314 J/(mol·K)) and T is the temperature in Kelvin. The correct answer is c.) 1.15 × 10^-9.
I hope this helps you with your problems! Let me know if you have any further questions.