Chapter 19

Define oxidation and reduction and explain the basic procedure for balancing redox reactions.

Explain the difference between a voltaic (or galvanic) electro-chemical cell and an electrolytic one.

Which reaction (oxidation or reduction) occurs at the anode of a voltaic cell? What is the sign of the anode? Do electrons flow toward or away from the anode?

Which reaction (oxidation or reduction) occurs at the cathode of a voltaic cell? What is the sign of the cathode? Do electrons flow toward or away from the cathode?

Explain the purpose of a salt bridge in an electrochemical cell.

Which unit is used to measure the magnitude of electrical current? Which unit is used to measure the magnitude of a potential difference? Explain how electrical current and potential difference differ.

What is the definition of the standard cell potential (Ecell)? What does a large positive standard cell potential imply about the spontaneity of the redox reaction occurring in the cell? What does a negative standard cell potential imply about the reaction?

Describe the basic features of a cell diagram (or line notation) for an electrochemical cell.

Why do some electrochemical cells employ inert electrodes such as platinum?

Describe the standard hydrogen electrode (SHE) and explain its use in determining standard electrode potentials.

How is the cell potential of an electrochemical cell (Ecell) related to the potentials of the half-cells?

Does a large positive electrode potential indicate a strong oxidizing agent or a strong reducing agent? What about a large negative electrode potential?

Explain why Ecell, Grxn, and K are all interrelated.

Does a redox reaction with a small equilibrium constant (K < 1) have a positive or a negative Ecell? Does it have a positive or a negative Grxn?

How does Ecell depend on the concentrations of the reactants and products in the redox reaction occurring in the cell? What effect does increasing the concentration of a reactant have on Ecell? Increasing the concentration of a product?

What are the anode and cathode reactions in a lead-acid storage battery? What happens when the battery is recharged?

What is a fuel cell? What is the most common type of fuel cell, and what reactions occur at its anode and cathode?

The anode of an electrolytic cell must be connected to which terminal-positive or negative-of the power source?

What is overvoltage in an electrochemical cell? Why is it important?

How is the amount of current flowing through an electrolytic cell related to the amount of product produced in the redox reaction?

What is corrosion? Why is corrosion only a problem for some metals (such as iron)?

Explain the role of each of the following in promoting corrosion: moisture, electrolytes, and acids.

How can the corrosion of iron be prevented?

Balance each redox reaction occurring in basic aqueous solution. a. H2O2(aq) + ClO2(aq)-ClO2-(aq) + O2(g) b. Al(s) + MnO4-(aq)-MnO2(s) + Al(OH)4-(aq) c. Cl2(g)-Cl-(aq) + ClO-(aq)

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. 2 Ag+(aq) + Pb(s)-2 Ag(s) + Pb2+(aq) b. 2 ClO2(g) + 2 I-(aq)-2 ClO2-(aq) + I2(s) c. O2(g) + 4 H+(aq) + 2 Zn(s)-2 H2O(l) + 2 Zn2+(aq)

Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. Ni2+(aq) + Mg(s)-Ni(s) + Mg2+(aq) b. 2 H+(aq) + Fe(s)-H2(g) + Fe2+(aq) c. 2 NO3-(aq) + 8 H+(aq) + 3 Cu(s)-2 NO(g) + 4 H2O(l) + 3 Cu2+(aq)

Determine whether or not each redox reaction occurs spontaneously in the forward direction. a. Ni(s) + Zn2+(aq)-Ni2+(aq) + Zn(s) b. Ni(s) + Pb2+(aq)-Ni2+(aq) + Pb(s) c. Al(s) + 3 Ag+(aq)-Al3+(aq) + 3 Ag(s) d. Pb(s) + Mn2+(aq)-Pb2+(aq) + Mn(s)

Use tabulated electrode potentials to calculate Grxn for each reaction at 25 C. a. Pb2+(aq) + Mg(s)-Pb(s) + Mg2+(aq) b. Br2(l) + 2 Cl-(aq)-2 Br-(aq) + Cl2(g) c. MnO2(s) + 4 H+(aq) + Cu(s)-Mn2+(aq) + 2 H2O(l) + Cu2+(aq)

Use tabulated electrode potentials to calculate Grxn for each reaction at 25 C. a. 2 Fe3+(aq) + 3 Sn(s)-2 Fe(s) + 3 Sn2+(aq) b. O2(g) + 2 H2O(l) + 2 Cu(s)-4 OH-(aq) + 2 Cu2+(aq) c. Br2(l) + 2 I-(aq)-2 Br-(aq) + I2(s)

Make a sketch of a concentration cell employing two Zn>Zn2+ halfcells. The concentration of Zn2+ in one of the half-cells is 2.0 M, and the concentration in the other half-cell is 1.0 * 10-3 M. Label the anode and the cathode and indicate the half-reaction occurring at each electrode. Also indicate the direction of electron flow.

Determine whether or not each metal, if coated onto iron, would prevent the corrosion of iron. a. Mg b. Cr c. Cu

Which products are obtained in the electrolysis of molten NaI?

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution: a. Ni(NO3)2(aq) b. KCI(aq) c. CuBr2(aq)

Make a sketch of an electrolytic cell that electroplates copper onto other metal surfaces. Label the anode and the cathode and indicate the reactions that occur at each.

Make a sketch of an electrolytic cell that electroplates nickel onto other metal surfaces. Label the anode and the cathode and indicate the reactions that occur at each.

What voltage can theoretically be achieved in a battery in which lithium metal is oxidized and fluorine gas is reduced? Why might such a battery be difficult to produce?

A metal forms the fluoride MF3. Electrolysis of the molten fluoride by a current of 3.86 A for 16.2 minutes deposits 1.25 g of the metal. Calculate the molar mass of the metal.

A sample of impure tin of mass 0.535 g is dissolved in strong acid to give a solution of Sn2+. The solution is then titrated with a 0.0448 M solution of NO3 -, which is reduced to NO(g). The equivalence point is reached upon the addition of 0.0344 L of the NO3- solution. Find the percent by mass of tin in the original sample, assuming that it contains no other reducing agents.

A 0.0251-L sample of a solution of Cu+ requires 0.0322 L of 0.129M KMnO4 solution to reach the equivalence point. The products of the reaction are Cu2+ and Mn2+. What is the concentration of the Cu2+ solution?

Which oxidizing agent oxidizes Br- but not Cl-? a. K2Cr2O7 (in acid) b. KMnO4 (in acid) c. HNO3