V. Everything’s Related Positive E° cell means spontaneous. Negative ΔG° means spontaneous. K >...
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![Page 1: V. Everything’s Related Positive E° cell means spontaneous. Negative ΔG° means spontaneous. K > 0 means spontaneous. Thus, all of these must be related.](https://reader035.fdocument.pub/reader035/viewer/2022062309/56649e385503460f94b2851e/html5/thumbnails/1.jpg)
V. Everything’s RelatedV. Everything’s Related
• Positive E°cell means spontaneous.
• Negative ΔG° means spontaneous.• K > 0 means spontaneous.• Thus, all of these must be related somehow.
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V. Potential and WorkV. Potential and Work
• Potential difference can be expressed as a function of work.
• We will use this new view to derive the relationship between E°cell and ΔG°.
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V. Eqn. Relating E°V. Eqn. Relating E°cellcell and ΔG° and ΔG°
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V. Sample ProblemV. Sample Problem
• Using tabulated half-cell potentials, calculate ΔG° for the reaction 2Na(s) + 2H2O(l) H2(g) + 2OH-
(aq) + 2Na+(aq).
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V. Potential and Equilibrium V. Potential and Equilibrium ConstantsConstants
• Using our new equation that relates standard cell potential to standard free energy, we can derive an equation between E°cell and K.
• We start with the equation that relates ΔG° to K.
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V. Eqn. Relating E°V. Eqn. Relating E°cellcell and K and K
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V. Sample ProblemV. Sample Problem
• Use tabulated half-cell potentials to calculate the equilibrium constant for the two electron oxidation of copper metal by H+.
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V. Relating ΔG°, K, and E°V. Relating ΔG°, K, and E°cellcell
• With the last equation derived, we summarize how we can convert between ΔG°, K, and E°cell.
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V. Nonstandard PotentialsV. Nonstandard Potentials• Just like any other system, electrochemical
cells may not be under standard conditions.
Why is the potential higher?
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V. Calculating Nonstandard V. Calculating Nonstandard Cell PotentialsCell Potentials
• Although we can qualitatively predict whether Ecell is higher or lower than E°cell, we’d like to calculate an exact value.
• We can derive an equation starting with the nonstandard ΔG equation.
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V. Deriving the Nernst EquationV. Deriving the Nernst Equation
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V. Nernst Eqn. GeneralizationsV. Nernst Eqn. Generalizations
• Under standard conditions, log Q = log 1 = 0, so Ecell will equal E°cell.
• If Q < 1, there are more reactants than products, so redox reaction shifts right; Ecell will be > than E°cell.
• If Q > 1, there are more products than reactants, so redox reaction shifts left; Ecell will be < than E°cell.
• If Q = K, then Ecell = 0.
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V. Sample ProblemV. Sample Problem
• Calculate the cell potential for the electrochemical cell represented by Ni(s)|Ni2+(aq, 2.0 M)||VO2
+(aq, 0.010 M), H+
(aq, 1.0 M), VO2+(aq, 2.0 M)|Pt(s).
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V. Concentration CellsV. Concentration Cells
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V. Concentration CellsV. Concentration Cells
• If the two half-cells are the same, there is no reason to reduce something on one side and oxidize the same thing on the other side.
• However, if there are [ ] differences, there is a push to get to equilibrium.
• Electrons flow in order to increase the [ ] of the dilute cell and decrease the [ ] of the concentrated cell.
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VI. Using ElectrochemistryVI. Using Electrochemistry
• If constructed correctly, electrochemical cells can be used to store and deliver electricity.
• We briefly look at dry-cell batteries, lithium ion batteries, and fuel cells.
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VI. Dry-cell BatteriesVI. Dry-cell Batteries
• Called dry-cell because there’s very little water.
• Voltage derived from the oxidation of Zn(s) and the reduction of MnO2(s).
• In acidic battery, MnO2(s) reduced to Mn2O3(s).
• In alkaline battery, MnO2(s) reduced to MnO(OH)(s).
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VI. Lithium Ion BatteriesVI. Lithium Ion Batteries
• This rechargeable battery works a bit differently.
• Motion of Li+ from anode to cathode causes e-’s to flow externally and reduce the transition metal.
• Battery is recharged by oxidizing the transition metal in the cathode.
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VI. Fuel CellsVI. Fuel Cells
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VII. Nonspontaneous Redox VII. Nonspontaneous Redox ReactionsReactions
• Nonspontaneous reactions can be forced to occur by inputting energy.
• Nonspontaneous redox reactions can be forced to occur by inputting energy in the form of electrical current.
• This type of electrochemical cell is called an electrolytic cell.
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VII. Voltaic vs. Electrolytic CellsVII. Voltaic vs. Electrolytic Cells
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VII. Predicting ProductsVII. Predicting Products
• In an aqueous electrolytic cell, it’s possible that H2O can be reduced or oxidized.
• To predict products, must consider potentials of all processes that could occur.
• Processes that are easiest (least negative or most positive half-cell potential) will occur.
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VII. OvervoltageVII. Overvoltage
• There is one problem in predicting electrolysis products.
• Some half-cell reactions don’t occur at their expected voltage potentials!
• e.g. 2H2O(l) O2(g) + 4H+(aq) + 4e- has
Eox = -0.82 V when [H+] = 1 x 10-7 M. However, kinetic factors require a voltage
of 1.4 V for this to occur.
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VII. Example ElectrolysisVII. Example Electrolysis
• What happens when a solution of NaI undergoes electrolysis?
• Oxidation 2I-
(aq) I2(aq) + 2e- E°ox = -0.54 V
2H2O(l) O2(g) + 4H+(aq) + 4e- E°ox = -1.4 V
• Reduction Na+
(aq) + e- Na(s) E°red = -2.71 V
2H2O(l) + 2e- H2(g) + 2OH-(aq) E°red = -0.41 V
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VII. Electrolysis of NaI SolutionVII. Electrolysis of NaI Solution
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VII. Electrolysis StoichiometryVII. Electrolysis Stoichiometry
• The # of electrons in any redox reaction can be used as a stoichiometric ratio.
• If we know how long an electrolysis takes place, and the magnitude of current that flowed, we can do stoichiometric calculations.
• Important relationships: 1 A = 1 C/s F = 96,485 C/mole e-
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VII. Sample ProblemVII. Sample Problem
• Copper can be plated out of a solution containing Cu2+ according to the half-reaction: Cu2+
(aq) + 2e- Cu(s). How long (in minutes) will it take to plate 10.0 g of copper using a current of 2.0 A?