A. Oxidation state or oxidation number is the “apparent” charge on an atom in a molecule or an ion.
B. Oxidation is a process that removes electrons from an atom, molecule, or ion.
C. Reduction is a process that adds electrons to an atom, molecule, or ion.
D. Oxidation-reduction reactions, or redox reactions, are those in which one or more species undergoes oxidation, and one or more other species undergoes reduction.
E. The field of chemistry that studies the relationship between electricity and chemical reactions is called electrochemistry.
A. We determine whether a reaction is a redox reaction by monitoring the oxidation numbers of the elements involved.
B. Because of conservation of electrons, oxidation and reduction must occur simultaneously. Further, the number of electrons removed by oxidation must equal the number of electrons added by reduction.
C. The substance being oxidized is called a reducing agent, or reductant.
D. The substance being reduced is called an oxidizing agent, or oxidant.
A. In balancing redox reactions, the number of atoms of each element must be the same on both sides of the equation. Additionally, the total charge must be the same on both sides of the equation.
B. The oxidation number method is summarized.
C. The method of half-reactions, or the ion-electron method, is summarized for acidic and basic media.

A. Voltaic cells, or galvanic cells, are devices for performing a spontaneous electrochemical reaction in which the electron transfer is forced to occur through an external pathway.
B. In many voltaic cells, the two half-reactions are contained in separate vessels; each vessel and its contents are referred to as half-cells. The two half-cells are connected by (1) an external circuit between the electrodes, and (2) a salt bridge between the two solutions.
C. The electrode at which oxidation occurs is called the anode. It has a negative charge in a voltaic cell.
D. The electrode at which reduction occurs is called the cathode. It has a positive charge in a voltaic cell.
E. The salt bridge creates a path for ion migration without mixing the contents of the two half-cells.
A. In spontaneous electrochemical reactions, the electrons flow from anode to cathode because it is energetically favorable for them to do so. The potential difference of an electron in the anode and an electron in the cathode is measured in volts.
B. The potential energy difference is the driving force of a voltaic reaction. The voltage is commonly referred to as the electromotive force, emf, or cell potential.
C. Under standard conditions of 25 o C, 1 atm pressure, and 1 M solutions, the voltage is referred to as the standard emf or the standard cell potential, E o cell.
D. The cell potential can be thought of as the combination of half-cell potentials.
1. The half-cell potential due to loss of electrons
at the anode is the oxidation potential, E ox
.
2. The half-cell potential due to electron gain at
the cathode is the reduction potential, E red
.
3. The cell potential, E cell
, is the sum of the half-
cell potentials.
4. The standard emf is the sum of the standard oxidation
and the standard reduction potentials.
E. Oxidation or reduction potentials cannot be directly measured for an isolated half-reaction. The reference half-reaction is the reduction of hydrogen

ion, 2 H + (aq) (1 M) + 2 e <---> H
2
(g) (1 atm), for which E ored = E oox = 0.00 V at 25 o C.
F. The half-cell potential for a reduction is equal in magnitude but opposite in sign to the same half-reaction written as an oxidation.
G. We tabulate only the standard reduction potentials or standard electrode potentials for half-reactions.
H. The more positive the E o for a half-reaction, the greater the tendency for the reaction to occur as written. The half-reactions with the most positive reduction potentials are the most powerful oxidizing agents, and those with the most positive oxidation potentials are the most powerful reducing agents.
A. A positive emf indicates a spontaneous process, and a negative emf indicates a nonspontaneous process.
B. The free energy change equals -nFE, where n is the number of moles of electrons transferred in the
reaction and F is Faraday‘s constant.
C. The equilibrium constant K at 298 K is given by the equation: E o = [(0.0591 V)/n] log K
D. Cell emf varies with the concentrations of the participating species. At 298 K, the Nernst equation is E = E o - [(0.0591 V)/n] log Q.
A. A battery is a voltaic cell that can consist of a single cell or several cells connected in series.
B. The lead storage battery, dry cell, and nickel- cadmium battery are described in detail.
C. Fuel cells are batteries in which conventional fuels, such as H
2
or CH
4
, are used.
Discussion Question: Discuss the various sizes and shapes of batteries. Does the physical size of the battery affect, within reason, the voltage output of the battery?
A. An electrochemical reaction that is nonspontaneous can be driven to proceed by applying an external electrical supply. Such reactions are called electrolysis reactions and are performed in electrolytic cells. Oxidation occurs at the anode,

and reduction occurs at the cathode. In an electrolytic cell, the anode is positive, and the cathode is negative.
B. Electrolysis of molten salts generally requires high temperatures to melt the salts.
C. The electrolysis of various aqueous solutions is discussed in detail.
D. In order for electrolysis to occur, the applied emf must be greater than, or equal to, the standard emf.
Because of internal resistance and the overvoltage phenomenon, the actual applied emf is invariably greater than the standard emf.
E. Active electrodes do participate in electrolysis, either by being oxidized or by being reduced.
Discussion Question: There are many commercial applications of electrolysis. Can you or the students propose other materials that could be produced by electrolysis? Is the popular technique of hair electrolysis consistent with the students’ understanding of electrolysis?
A. Stoichiometrically, the number of ions of a metal that are reduced by an applied electrical current is proportional to the number of electrons delivered to the cell by the current.
B. A coulomb is the amount of charge passing a point in a circuit operating at 1 ampere for 1 second.
Therefore, coulombs = (amperes) (seconds).
C. The amount of charge passed through a circuit is related through the Faraday constant to the number of moles of electrons passed.
D. The maximum work that can be derived from a galvanic cell is -nFE.
Discussion Question: Discuss the amount of electricity used by a plant that produces aluminum electrolytically. Propose an amount of aluminum to be produced in a day, and apply your local electric rates.
A. Corrosion is a redox reaction in which a metal is attacked by its environment to form an unwanted substance.
B. The corrosion of iron, commonly referred to as rusting, requires the presence of oxygen and water and is discussed in detail.

Discussion Question: Modern materials, like plastics and ceramics, do not corrode. Why haven’t these materials replaced structural metals? What typical applications of metals have been taken over by plastics and ceramics?
1. Define oxidation and reduction.
2. Define reducing agent and oxidizing agent.
3. Given the reaction ClO3-(aq) + 3 S2-(aq) + 3 H2O(l)-->
Cl-(aq) + 3 S(s)+ 6 OH-(aq). Which substance is the reducing agent? Which substance is the oxidizing agent?
4. Balance the following redox reactions in acidic solution:
(a) Fe3+(aq) + I-(aq) --> Fe2+(aq) + I2(s);
(b) NO3-(aq)+ H+(aq)+ I2(s) --> NO2(g)+ H2O(l)+ IO3-(aq);
© H2O2(aq)+ Cr2O72-(aq) --> Cr3+(aq) + O2(g).
5. Balance the following redox reactions in basic solution:
(a) Cl2(g) --> Cl-(aq) + ClO-(aq);
(b) CN-(aq) + MnO4-(aq) --> MnO2(s) + CNO-(aq).
6. Complete and balance each of the following: (a) NCl3(g) +
H2O(l) --> NH3(g) + HOCl(aq); (b) FeO(s) + O2(g) -->
Fe3O4(s); © PbO2(s) + N2H4(g) --> Pb(s) + N2(g) + H2O(g).
7. The standard potentials for reducing the following ions are: Mg2+,-2.37 V; Fe2+,-0.44 V; Cu2+, 0.34 V; Ag+, 0.80 V.
(a) Which of the following is the best reducing agent: Mg,
Cu, Ag+, or Mg2+? Explain briefly. (b) What is the emf of a standard cell utilizing the following reaction: 2 Ag+ +
Fe --> 2 Ag + Fe2+? © Will the following reaction be spontaneous: Mg + Cu2+ --> Mg2+ + Cu? Explain briefly.
8. Consider the following electrode potentials: Fe2+ + 2e- -->
Fe, Eo = -0.44 V, and Cu2+ + 2e- --> Cu, Eo = 0.34 V.
(a) What voltage would be produced by a standard voltaic cell that utilizes these half-reactions? (b) Draw the

cell, labeling the anode, the cathode, and the directions of the ion and electron motions. © What effect will increasing the concentration of Cu2+ have on the cell emf?
9. The reaction Zn + Cu2+ --> Cu + Zn2+ produces a standard potential of 1.10 V. What voltage is produced when [Cu2+]=
0.50 M and [Zn2+] = 0.0050 M?
10. Eo = 0.76 V for Zn + 2 H+ <--> Zn2+ + H2. What is the free energy change and equilibrium constant for this reaction?
11. What is the difference between a voltaic cell and an electrolytic cell?
12. If a current of 6.00 A is passed through a bath of aqueous
CuSO4 for 2.00 hrs, how much copper metal would plate out?
13. What is the minimum energy required to produce 1.0 mol of
Cu by electrolysis if the process requires an emf of 1.0 V and is 100% efficient?
14. Describe and explain the relative merits of paint, tin, and zinc as protective coatings on iron.
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