Chapter 4 - Oxidation-reduction

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Oxidation and Reduction (Ch. 4)
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 diagrammatic tools that help us to rationalize and predict redox behavior, reactivity, and stability.
Learning goals:
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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 nonredox reactions, reactions that are and are not pH-dependent, and ultimately to predict and
rationalize stability, reactivity, corrosion, and passivation.
Introduction
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.
Balancing Redox Reactions
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:
Cr2O72- + I- = Cr3+ + IO3-
(balance in acid)
MnO4- + HCHO = MnO2 + CO32- (balance in base)
Electrochemical reactions
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 Q
E = E° – (0.0592/n) log Q
2 H+ + 2 e–  H2
What is E°1/2 for this reaction?
What is E°1/2 for this reaction at pH 5 and PH2 = 1 atm?
(And … what is wrong with this question?)
2 H+ + 2 e–  H2 and O2 + 4 H+ + 4 e–  2 H2O
What is the difference between the H2/H+ and O2/H2O couples at pH 5?
Stability field of water
Redox stability diagrams
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:
Latimer diagrams
What is E° for MnO4–  Mn2+?
What is E° for MnO4–  MnO2?
Latimer diagrams
Unstable species have a lower number to the left than to the right
Latimer diagrams
Which Mn species are unstable with respect to disproportionation?
Latimer diagrams
Let’s take a closer look at the stability of MnO42–
Practical considerations
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 (NO2, NO, N2H4) are
unstable relative to the elements (O2, N2, and H2).
Identifying stable and unstable oxidation states is
easy using a Frost diagram
Frost diagrams
What is a Frost diagram?
How do we define an element on a Frost diagram?
Frost diagram for Mn
Frost diagrams
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
Pourbaix diagram
Plot of electrochemical equilibria as a function of pH
Key equilibria for Fe system – which are (are not) pH-dependent?
Fe2+ + 2 e–  Fe(s)
Fe3+ + e–  Fe2+
Fe3+ + 3 OH–  Fe(OH)3(s)
Fe2+ + 2 OH–  Fe(OH)2(s)
Fe(OH)3 + e– + 3 H+  Fe2+ + 3 H2O
Pourbaix diagram
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 for Fe (simplified)
• What do the lines mean?
• What is the slope of the line between Fe2+(aq) and Fe2O3(s)
Pourbaix diagram for Fe
• What can we say about the stability of Fe(s) in H2O?
• Under what conditions is Fe passivated or protected against corrosion?

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