Unit 9: Thermodynamics and Electrochemistry.

Justify the sign of entropy or relative entropy values by examining the dispersal of states (number of microstates) and number of moles of gas during a change.

1. Does the phase change H₂O(l) → H₂O(g) results in an increase of decrease in entropy?
2. What happens to entropy at the particle level during the reaction: H₂(g) + O₂(g) → H₂O(l) occurs?

Justify the impact of temperature on entropy.

1. Explain why the entropy of a sample of argon gas increases when its temperature is increased from 25°C to 50°C.
2. Two samples of nitrogen gas are held at the same pressure. Sample A is at 25°C and Sample B is at 100°C. Which sample has the greater entropy? Justify your answer.

Calculate the entropy of a system given the entropies of component species using ∆S°rxn = ∑∆S°prod - ∑∆S°react.

1. Calculate ΔS°rxn for the reaction: N2(g) + 3H2(g) → 2NH3(g) Given: S°(N2) = 192.8 J mol^-1 K^-1 S°(H2) = 130.7 J mol^-1 K^-1 S°(NH3) = 192.5 J mol^-1 K^-1

Use the given value of ∆G to determine if a system is thermodynamically favorable at a certain temperature.

1. A reaction has ΔG = -35 kJ mol^-1 at 298 K. Is the reaction thermodynamically favorable? Explain.
2. A reaction has ΔG = +12 kJ mol^-1 at 298 K. What does this indicate about the spontaneity of the reaction under these conditions?

Calculate the free energy of a system using ∆G°rxn = ∑∆Gf°prod - ∑∆Gf°react.

1. Calculate ΔG°rxn for the reaction: H₂(g) + Cl₂(g) → 2HCl(g) Given: ΔGf°(HCl) = -95.3 kJ mol^-1 ΔGf°(H₂) = 0 kJ mol^-1 ΔGf°(Cl₂) = 0 kJ mol^-1

Determine thermodynamic favorability using ∆G = ∆H - T∆S at a given temperature.

1. Determine ΔG at 298 K for a reaction with ΔH = -80 kJ mol^-1 and ΔS = -150 J mol^-1 K^-1. Is the reaction thermodynamically favorable?

Justify (∆G) thermodynamic favorability based on the signs of ∆H and ∆S at high and low temperatures.

1. For each combination below, determine whether the reaction is favorable at all temperatures, no temperatures, only high temperatures, or only low temperatures: (a) ΔH < 0, ΔS > 0, (b) ΔH > 0, ΔS < 0, (c) ΔH < 0, ΔS < 0, (d) ΔH > 0, ΔS > 0

Identify reasons why a thermodynamically favorable reaction might proceed at such a slow rate that it does not occur measurably.

1. Carbon and graphite and both allotropes of Carbon (C), meaning they exist in different forms in the same state. The reaction: C_(graphite) → C_(diamond) has a ∆G° of -2.9kJ mol^-1. However graphite does not convert to diamond at any measureable rate. Suggest an explanation for these 2 facts.

Contrast thermodynamically favorable and kinetically controlled reactions.

1. Explain why the combustion of gasoline is thermodynamically favorable but does not occur spontaneously at room temperature.

Relate K to ∆G° using ∆G° = -RT ln K or it's rearranged equation K = e^-∆G°/RT

1. The reaction A ⇌ B has K_eq=1000 at 25^oC. What is the value of ∆G? Is this reaction thermodynamically favourable at this temperature?
2. The reaction A ⇌ B has ΔG° = −125 J/mol at 298 K. Determine the value of the equilibrium constant (K_eq) at this temperature.

Use a stepwise analysis of dissolution to determine if a salt is likely to dissolve based on enthalpy and entropy considerations.

1. What are the 3 attractive forces that influence dissolution. Is each endothermic or exothermic?

Describe how two reactions can be coupled to make a thermodynamically favorable process occur.

1. Explain why ATP hydrolysis is often coupled to nonspontaneous biological reactions.
2. A reaction has ΔG = +15 kJ mol^-1. It is coupled to a second reaction with ΔG = -30 kJ mol^-1. Determine the overall ΔG and whether the coupled process is thermodynamically favorable.

Identify the components of an electrochemical cell and the role each part plays.

1. Identify the function of the salt bridge in a galvanic cell.
2. Explain why a galvanic cell would stop operating if the salt bridge were removed.

Determine the direction of electron flow in an electrochemical cell.

Represent an electrochemical cell with an appropriate particle diagram.

1. Draw a particle-level diagram of a Zn/Cu galvanic cell showing the movement of electrons and ions.

Label the anode and cathode with the corresponding redox process.

Use signs of ∆G and/or E°cell to distinguish between galvanic and electrolytic cells.

1. Is a cell with a ∆G = -1000kJ/mol thermodynamicaly favourable?
2. A electrochemical cell is measured to have an electrochemical potential (E_cell) of 1.20V. Is the cell thermodynamically favourable or thermodynamically unfavourable?

Calculate the E°cell of an electrochemical cell.

1. Calculate E°cell for a galvanic cell composed of: Cu2+ + 2e^- → Cu E° = +0.34 V Zn2+ + 2e^- → Zn E° = -0.76 V

Apply the equation ∆G = -nFE°cell.

1. Calculate ΔG° for a reaction with n = 2 and E°cell = 1.10 V. Use F = 96485 C mol^-

Explain how a concentration cell works.

1. A concentration cell is constructed using two Cu/Cu2+ half-cells. One contains 1.0 M Cu2+ and the other contains 0.010 M Cu2+. Explain why a voltage is produced.

Understand why Le Châtelier’s principle cannot be used to explain electrochemical systems.

1. Explain why Le Châtelier's principle cannot be directly applied to predict changes in cell potential.

Use comparisons of Q values to determine how changes in the cell as it progresses impact E relative to E°.

1. For a galvanic cell, predict whether E is greater than, less than, or equal to E° when Q < 1.

Apply the Nernst equation qualitatively to discuss the impact of concentration on cell potential.

1. Predict how increasing the concentration of Cu2+ in a Cu/Zn galvanic cell would affect the cell potential.

Apply stoichiometric calculations using Faraday’s Law (I = q/t) to determine: 1. Number of electrons transferred, 2. change in mass of electrodes, 3. current, time, and charge of ionic species.

1. A current of 2.50 A is passed through an electrolytic cell for 30.0 minutes. Calculate the total charge transferred.
2. How many moles of electrons are transferred when a charge of 9650 C passes through a cell?
3. Silver ions are reduced according to: Ag+ + e^- → Ag(s) Calculate the mass of silver deposited when 0.250 mol of electrons pass through the cell.