Chemistry Outcomes Review: Chapter 10

Reduction-Oxidation: Electrochemistry

In this chapter we introduced two kinds of electrochemical cells: electrolytic and galvanic. In electrolytic cells, we use an electric current to produce chemical reactions by reducing and oxidizing species in solution. In galvanic cells oxidation and reduction half reactions are spatially separated and the free energy of a chemical reaction drives electrons through an external circuit where they can be used to do work.

The quantitative redox relationships that follow from electrochemical cell measurements also characterize redox reactions that occur between reactants mixed in the same solution. Redox reactions in living cells obey these same relationships, which can be used to help understand and explain observed stoichiometric relationships in organisms.

Check your understanding of the ideas in the chapter by reviewing these expected outcomes of your study.

You should be able to:

• Use evidence from experimental observations to write probable half reactions for the reductions and oxidations taking place in an electrolytic cell [Section 10.1].
• Describe and draw molecular level diagrams of the processes going on and the flow of charge in an electrochemical cell [Sections 10.1 and 10.2].
• Use the Faraday and relationships among time, electric current, cell potential, and cell reaction stoichiometry to calculate the amounts of products from an electrolysis or the amount of work available from the reactants in a galvanic cell [Sections 10.1 and 10.3].
• Show how to connect two metal-metal ion half cells to make a galvanic cell and explain the role of each component of the cell [Section 10.2].
• Use the known cell reaction to identify the anode and cathode of an electrochemical cell [Sections 10.1, 10.2, 10.3, and 10.4].
• Use the known sign for a galvanic cell potential to identify the direction of the cell reaction and the anode and cathode of the cell [Section 10.3].
• Describe how an electrode senses the redox half reaction in a half cell in which the reduced and oxidized species are both present as dissolved ions and/or molecules [Section 10.4].
• Translate a physical cell set up to the conventional line notation for cells and vice versa [Section 10.4].
• Apply Le Chatelier's principle to predict the direction of change of cell potentials as concentrations in the half cells are changed [Sections 10.4, 10.5, 10.7, and 10.8].
• Determine free energy change for a cell reaction from cell potential and vice versa [Section 10.5].
• Use the Nernst equation, which relates the cell potential, the standard cell potential, and the reaction quotient for the cell reaction, to determine any one of these quantities, if the other two are known [Sections 10.5, 10.6, and 10.9].
• Use the standard cell potential to determine the equilibrium constant for a cell reaction and vice versa [Sections 10.5 and 10.8].
• Determine an unknown cell potential for a cell reaction by combining cell reactions with known cell potentials to give the desired cell reaction and its cell potential [Section 10.6].
• Use a table of standard reduction potentials to predict the direction of any redox reaction (for which data are given) and determine its standard cell potential [Sections 10.7, 10.8, 10.9, and 10.10].
• Use the Nernst equation to convert standard reduction potentials to reduction potentials under non-standard conditions and vice versa [Sections 10.8, 10.9, and 10.10].
• Use a table of standard reduction potentials at pH 7 to predict the direction and standard cell potential at pH 7 for redox reactions in biological systems [Sections 10.8, 10.9, and 10.10].
• Combine free energy changes for uncomplicated redox reactions with free energies for complexation and solubility equilibria in a redox system to get the free energy change and reduction potential for the actual net half reaction that occurs [Sections 10.9 and 10.10].
• Convert a stepwise series of reactions to a diagram that shows the coupling of reactions and vice versa [Section 10.10].
• Use standard reduction potentials to determine the probable sequence of a series of coupled redox reactions and vice versa [Section 10.10].