1. Big-picture perspective: Oxidation-reduction reactions are integral to many aspects of inorganic chemistry. Building on your existing knowledge of electrochemistry, we will discuss some fundamental aspects of inorganic electrochemistry that you may not have previously considered (and therefore make some new connections to other areas of chemistry) and also introduce and use three diagram- matic tools that help us to rationalize and predict redox behavior, reactivity, and stability. Learning goals: Balance complex oxidation-reduction reactions by the ion-electron method. Understand periodic trends in the activity series and electrochemical series. Use the Nernst equation to determine half-cell and cell potentials. Derive the stability field of water and use this to rationalize aqueous redox chemistry. Construct and be proficient with Latimer diagrams, using them to determine unknown reduction potential values and to quickly identify stable and unstable species. Construct and be proficient with Frost diagrams, using them to identify stable and unstable species, as well as those that are strong oxidizers. Construct and be proficient with Pourbaix diagrams, using them to identify redox and non- redox reactions, reactions that are and are not pH-dependent, and ultimately to predict and rationalize stability, reactivity, corrosion, and passivation. Oxidation and Reduction (Ch. 4)

2. Oxidation-reduction phenomena are integral to many aspects of inorganic chemistry Many elements, including the transition metals, have multiple accessible oxidation states. The compounds that they form, as well as their chemical properties and reactivity, are tied intimately to their oxidation states Many inorganic compounds catalyze, and participate in, redox reactions (e.g. in industry and biology) Energy conversion processes (solar, batteries, fuel cells) rely on inorganic redox reactions Review and know: Electrochemistry chapter in (any) general chemistry textbook – assigning oxidation states, balancing redox reactions, using and applying the table of standard reduction potentials, Nernst equation (important!), and quantitative relationships among E, G, and K (important!) Our focus in Chem 310: (a) thermochemical aspects of reduction potentials and their relationships to redox trends among the elements and (b) diagrammatic tools to help predict and rationalize electrochemical reactions. Introduction

3. The Ion-Electron Method: (1) Identify the elements undergoing redox, balance them in half rxns (2) Add water to balance O (3) Add H + to balance H (4) Add e - to balance charge (5) Combine half-reactions (6) For reactions in base, add OH - to neutralize H + Practice balancing: Cr 2 O 7 2- + I - = Cr 3+ + IO 3 - (balance in acid) MnO 4 - + HCHO = MnO 2 + CO 3 2- (balance in base) Balancing Redox Reactions

4. Let’s now take a look at some specific electrochemical reactions, emphasizing redox stability (e.g. conditions under which certain species are / are not stable, and/or under which certain redox reactions are / are not spontaneous.) In the process, we will investigate three diagrammatic tools, which are familiar to and used by practicioners of inorganic chemistry – Latimer diagrams, Frost diagrams, and Pourbaix diagrams. These types of diagrams conveniently and concisely compile and present electrochemical data, but each one has a unique “twist” and therefore each serves a distinct purpose in helping us to rationalize and predict electrochemical stability and reactivity. Recall several useful equations: ΔG° = – nFE°E = E° – (RT/nF) ln QE = E° – (0.0592/n) log Q Electrochemical reactions

5. What is E° 1/2 for this reaction? What is E° 1/2 for this reaction at pH 5 and P H2 = 1 atm? (And … what is wrong with this question?) 2 H + + 2 e –  H 2

6. What is the difference between the H 2 /H + and O 2 /H 2 O couples at pH 5? 2 H + + 2 e –  H 2 and O 2 + 4 H + + 4 e –  2 H 2 O

7. Stability field of water

8. Three types of redox stability diagrams are helpful for presenting similar information in ways that are useful in different situations Latimer diagrams – E°’s for successive redox reactions For Mn species in acid: Redox stability diagrams

9. What is E° for MnO 4 –  Mn 2+ ? Latimer diagrams What is E° for MnO 4 –  MnO 2 ?

10. Unstable species have a lower number to the left than to the right Latimer diagrams

11. Which Mn species are unstable with respect to disproportionation? Latimer diagrams

12. Let’s take a closer look at the stability of MnO 4 2– Latimer diagrams

13. We need to consider all possible disproportionation reactions We need to consider kinetics: thermodynamically unstable species can be quite stable kinetically. Most N-containing molecules (NO 2, NO, N 2 H 4 ) are unstable relative to the elements (O 2, N 2, and H 2 ). Identifying stable and unstable oxidation states is easy using a Frost diagram Practical considerations

14. What is a Frost diagram? How do we define an element on a Frost diagram? Frost diagrams

15. Frost diagram for Mn

16. Same information as in a Latimer diagram, but graphically shows stability and oxidizing power. Unstable species are above the lines connecting neighbors. Lowest species on the diagram are the most stable Highest species on the diagram are the strongest oxidizers Frost diagrams

17. Plot of electrochemical equilibria as a function of pH Key equilibria for Fe system – which are (are not) pH-dependent? Fe 2+ + 2 e –  Fe(s) Fe 3+ + e –  Fe 2+ Fe 3+ + 3 OH –  Fe(OH) 3 (s) Fe 2+ + 2 OH –  Fe(OH) 2 (s) Fe(OH) 3 + e – + 3 H +  Fe 2+ + 3 H 2 O Pourbaix diagram

18. Pourbaix diagram: Plot of E vs. pH How are pure redox reactions plotted? How are pure acid-base reactions plotted? How are “mixed” reactions plotted? Pourbaix diagram

19. Pourbaix diagram for Fe (simplified) What do the lines mean? What is the slope of the line between Fe 2+ (aq) and Fe 2 O 3 (s)

20. What can we say about the stability of Fe(s) in H 2 O? Under what conditions is Fe passivated or protected against corrosion? Pourbaix diagram for Fe