Chapter 20

In electrochemistry, spontaneous redox reactions are used for what purpose?

In electrochemistry, what kind of reaction can be driven by electricity?

Give the basic definitions of oxidation and reduction and explain the basic procedure for balancing redox reactions.

Explain the difference between a voltaic (or galvanic) electrochemical cell and an electrolytic cell.

What 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?

What 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.

What unit is used to measure the magnitude of electrical current? What 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 ( \(E_{\text {cell }}^{\circ}\) )? 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 ( \(E_{\text {cell }}^{\circ}\) ) 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 \(E_{\mathrm{cell}}^{\circ}, \Delta G_{\mathrm{rxn}}^{\circ},\) and \(K\) are all interrelated.

Use the Nernst equation to show that \(E_{\text {cell }}=E_{\text {cell }}^{\circ}\) under standard conditions.

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

Explain how a fuel-cell breathalyzer works.

List some applications of electrolysis.

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

Why does the electrolysis of an aqueous sodium chloride solution produce hydrogen gas at the cathode?

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?

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

Balance each redox reaction occurring in acidic aqueous solution. MISSED THIS? Read Section \(20.2 ;\) Watch \(\mathrm{KCV} 20.2,\) WE 20.2 a. \(\mathrm{K}(s)+\mathrm{Cr}^{3+}(a q) \longrightarrow \mathrm{Cr}(s)+\mathrm{K}^{+}(a q)\) b. \(\mathrm{Al}(s)+\mathrm{Fe}^{2+}(a q) \longrightarrow \mathrm{Al}^{3+}(a q)+\mathrm{Fe}(s)\) c. \(\mathrm{BrO}_{3}^{-}(a q)+\mathrm{N}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{Br}^{-}(a q)+\mathrm{N}_{2}(g)\)

Balance each redox reaction occurring in acidic aqueous solution. MISSED THIS? Read Section 20.2; Watch KCV 20.2, IWE 20.2 a. \(\mathrm{PbO}_{2}(s)+\mathrm{I}^{-}(a q) \longrightarrow \mathrm{Pb}^{2+}(a q)+\mathrm{I}_{2}(s)\) b. \(\mathrm{SO}_{3}^{2-}(a q)+\mathrm{MnO}_{4}^{-}(a q) \longrightarrow \mathrm{SO}_{4}^{2-}(a q)+\mathrm{Mn}^{2+}(a q)\) c. $\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{SO}_{4}^{2-}(a q)+\mathrm{Cl}^{-}(a q)

Balance each redox reaction occuring in acidic aqueous solution. a. \(\mathrm{I}^{-}(a q)+\mathrm{NO}_{2}^{-}(a q) \longrightarrow \mathrm{I}_{2}(s)+\mathrm{NO}(g)\) b. \(\mathrm{ClO}_{4}^{-}(a q)+\mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{ClO}_{3}^{-}(a q)+\mathrm{Cl}_{2}(g)\) c. \(\mathrm{NO}_{3}^{-}(a q)+\mathrm{Sn}^{2+}(a q) \longrightarrow \mathrm{Sn}^{4+}(a q)+\mathrm{NO}(g)\)

Make a sketch of a concentration cell employing two \(\mathrm{Zn} / \mathrm{Zn}^{2+}\) halfcells. The concentration of \(\mathrm{Zn}^{2+}\) in one of the half-cells is \(2.0 \mathrm{M}\), and the concentration in the other half- cell is \(1.0 \times 10^{-3} \mathrm{M}\). Label the anode and the cathode and indicate the half-reaction occurring at each electrode. Also indicate the direction of electron flow.

A concentration cell consists of two \(\mathrm{Sn} / \mathrm{Sn}^{2+}\) half- cells. The cell has a potential of \(0.10 \mathrm{~V}\) at \(25^{\circ} \mathrm{C} .\) What is the ratio of the \(\mathrm{Sn}^{2+}\) concentrations in the two half-cells?

Determine whether or not each metal, if coated onto iron, would prevent the corrosion of iron. a. Zn b. \(\mathrm{Sn}\) c. Mn

Write equations for the half-reactions that occur in the electrolysis of molten potassium bromide.

What products are obtained in the electrolysis of molten NaI?

Write equations for the half-reactions that occur in the electrolysis of a mixture of molten potassium bromide and molten lithium bromide.

Write equations for the half-reactions that occur at the anode and cathode for the electrolysis of each aqueous solution. a. \(\operatorname{NaBr}(a q)\) b. \(\mathrm{PbI}_{2}(a q)\) c. \(\mathrm{Na}_{2} \mathrm{SO}_{4}(a q)\)

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

Consider the molecular view of an electrochemical cell involv- ing the overall reaction: $$ \mathrm{Zn}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Ni}(s) $$ Draw a similar sketch of the cell after it has generated a substantial amount of electrical current.

The molar mass of a metal (M) is \(50.9 \mathrm{~g} / \mathrm{mol} ;\) it forms a chloride of unknown composition. Electrolysis of a sample of the molten chloride with a current of 6.42 A for 23.6 minutes produces \(1.20 \mathrm{~g}\) of \(\mathrm{M}\) at the cathode. Determine the empirical formula of the chloride.

A metal forms the fluoride \(\mathrm{MF}_{3}\). Electrolysis of the molten fluoride by a current of 3.86 A for 16.2 minutes deposits \(1.25 \mathrm{~g}\) of the metal. Calculate the molar mass of the metal.

To what \(\mathrm{pH}\) should you adjust a standard hydrogen electrode to get an electrode potential of \(-0.122 \mathrm{~V}\) ? (Assume that the partial pressure of hydrogen gas remains at 1 atm.

An electrochemical cell has a positive standard cell potential but a negative cell potential. Which statement is true for the cell? a. \(K>1 ; Q>K\) b. \(K<1 ; Q>K\) c. \(K>1 ; Q

A redox reaction employed in an electrochemical cell has a negative \(\Delta G_{\mathrm{rxn}}^{\circ} .\) Which statement is true? a. \(E_{\text {cell }}^{\circ}\) is positive; \(K<1\) b. \(E_{\text {cell }}^{\circ}\) is positive; \(K>1\) c. \(E_{\text {cell }}^{\circ}\) is negative; \(K>1\) d. \(E_{\text {cell }}^{\circ}\) is negative; \(K<1\)