Unit 9: Thermodynamics and Electrochemistry.

9.1 Introduction to Entropy.

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.

9.2 Absolute Entropy and Entropy Change.

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

9.3 Gibbs Free Energy and Thermodynamic Favorability.

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

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

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

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

9.4 Thermodynamic and Kinetic Control.

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.

9.5 Free Energy and Equilibrium.

Quantitatively relate the sign of ∆G° to the magnitude of K.

Qualitatively relate K to ∆G° using K = e⁻∆G°/RT or ∆G° = -RT ln K.

9.6 Free Energy of Dissolution.

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

9.7 Coupled Reactions.

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

9.8 Galvanic (Voltaic) and Electrolytic Cells.

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

Determine the direction of electron flow in an electrochemical cell.

Represent an electrochemical cell with an appropriate particle diagram.

Label the anode and cathode with the corresponding redox process.

9.9 Cell Potential and Free Energy.

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

Calculate the E°cell of an electrochemical cell.

Apply the equation ∆G = -nFE°cell.

9.10 Cell Potential Under Nonstandard Conditions.

Explain how a concentration cell works.

  1. Calculate the molar mass of a chemical species.

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

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

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

9.11 Electrolysis and Faraday’s Law.

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.

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