Question 4: Enthalpy of Reaction (12 points)
a. A high school maintenance worker (who knew a lot about chemistry) had two chemicals for removing ice from sidewalks. The labels weren't on the containers any longer, but he knew he had a container of NaCl and a container of MgCl2.
He took a sample from each container to the chemistry lab. After looking up the enthalpies of formation of the chemicals, he calculated the ΔHreaction for dissolving each powder: NaCl(s) Na+(aq) + Cl(aq), and MgCl2(s) Mg2+ + 2Cl–(aq). He then put a sample of each powder in a coffee-cup calorimeter and added water.
When the sample from container A dissolved, the temperature of the solution decreased by 0.02°C. When the sample from container B dissolved, the temperature increased by 0.58°C. Which chemical was in each container? Use the table of enthalpies of formation to help you. Explain your reasoning. (4 points)
I would assume that the samples contained the same moles of NaCl and MgCl2. The dissolving of MgCl2 gives off more heat than the dissolving of NaCl, and so the sample which gives the larger temperature increase is MgCl2.

b. Use Hess's law and the following equations to calculate the ΔHreaction for the reaction CO(g) + 3H2(g) CH4(g) + H2O(g). Show your work. (4 points)
C(s) + O2(g) CO(g) ΔH = –110.5 kJ/mol
C(s) + 2H2(g) CH4(g) ΔH = –74.85 kJ/mol
2H2(g) + O2(g) 2H2O(g) ΔH = –483.66 kJ/mol
ΔHrx = ΔHf(products) - ΔHf(reactants)
ΔHrx = -74.85 kJ + (-483.66 kJ) - ( -110.5 kJ) = -448.0 kJ

c. Draw a potential energy diagram for the formation of methane: C(s) + 2H2(g) CH4(g), ΔH = –74.85 kJ/mol. Label the graph with "Reactants," "Products," "ΔHreaction," and "Ea." (4 points)

Question 5: Entropy (6 points)
a. Predict whether the entropy would increase or decrease for the following phase changes. (2 points)
i. H2O(g) H2O(l) (1 point)
Decrease
ii. H2O(s) H2O(l) (1 point)
Increase
b. Predict whether the entropy would increase or decrease in the reaction KCl(s) K+(aq) + Cl–(aq). Explain your answer. (2 points)
In the crystal of KCl the ions are relatively fixed, but when the ions go into solution they can move in many more positions. So the entropy increases in the direction of dissolving.

c. Briefly state the second law of thermodynamics. Explain how the second law of thermodynamics is upheld when an aqueous salt solution is evaporated to form salt crystals again. (2 points)
The second law states that the entropy of a closed system will always increase. Entropy is a measure of "order"; a system always moves towards a state of less order. The idea of entropy increasingly explains why a crystal of salt will dissolve in a beaker of water of its own accord.

Question 6: Gibbs Free Energy (8 points)
a. What is the Gibbs free energy at 298 K for the breakdown of ammonia?
2NH3(g) N2(g) + 3H2(g)
ΔH = 91.8 kJ/mol
ΔS = 198.7 J/(mol·K)
Is the reaction spontaneous at 298 K (room temperature)? (2 points)
ΔG = ΔH - TΔS
ΔG = 91.8 kJ - 298K(0.1987 kJ/K)
ΔG = 32.6 kJ
No, the reaction is not spontaneous at room temperature

b. What predictions can you make about ΔH, ΔS, and ΔG for the general reaction A(s) + 2B(g) 3C(g) + D(g) + heat? What predictions can you make about the spontaneity of the reaction? Explain your reasoning. (4 points)
Heat is a product so the reaction is exothermic, and ΔH < 0. ΔS > 0 because you are going from 2 moles of gaseous reactant to 4 moles of gaseous products. Therefore, ΔG will always be less than 0, and the reaction is always spontaneous.

c. The reaction 2NO2(g) N2(g) + 2O2(g) is spontaneous at all temperatures, yet very little of this toxic air pollutant dissociates under normal conditions. How can this be? (2 points)
The reaction has a very large activation energy.

Could someone please check my work?