Deck 19: Electrochemistry

Question 1

Explain a redox reaction with an example.

Answer 1

In any electrochemical process, electrons flow from one chemical substance to another, driven by an oxidation-reduction (redox) reaction. A redox reaction occurs when electrons are transferred from a substance that is oxidized to one that is being reduced. The reductant is the substance that loses electrons and is oxidized in the process the oxidant is the species that gains electrons and is reduced in the process. A redox reaction can be described as two half-reactions, one representing the oxidation process and one the reduction process. For the reaction of zinc with bromine, the overall chemical reaction is as follows:
Zn(s) + Br₂(aq) → Zn²⁺(aq) + 2Br⁻(aq)
The half-reactions are as follows:
Reduction half-reaction: Br₂(aq) + 2e⁻ → 2Br⁻(aq)
Oxidation half-reaction: Zn(s) → Zn²⁺(aq) + 2e⁻
Zn is the reductant in this reaction (it loses electrons), and Br₂ is the oxidant (it gains electrons).

Question 2

The anode is the site of the reduction half-reaction.

Answer 2

False

Question 3

What is the purpose of a salt bridge in a galvanic cell?

Answer 3

A salt bridge is a U-shaped tube inserted into both solutions of a galvanic cell that contains a concentrated liquid or gelled electrolyte and completes the circuit between the anode and the cathode. The ions in the salt bridge are selected so that they do not interfere with the electrochemical reaction by being oxidized or reduced themselves or by forming a precipitate or complex; commonly used cations and anions are Na⁺ or K⁺ and NO₃⁻ or SO₄²⁻, respectively. The electrolyte in the salt bridge serves two purposes: it completes the circuit by carrying electrical charge and maintains electrical neutrality in both solutions by allowing ions to migrate between them. The identity of the salt in a salt bridge is unimportant, as long as the component ions do not react or undergo a redox reaction under the operating conditions of the cell. In the absence of a salt bridge or some other similar connection, the reaction would rapidly cease because electrical neutrality could not be maintained.

Question 4

A galvanic cell differs from an electrolytic cell in that a galvanic cell _____.

A) consumes electrical energy from an external source
B) drives a nonspontaneous redox reaction
C) contains two electrodes to provide an external connection
D) uses the energy released during a spontaneous redox reaction to generate electricity
E) contains an electrolyte that connects the electrodes

Question 5

A voltaic cell consumes electrical energy from an external source.

Question 6

The oxidant is the species that loses electrons in the process.

Question 7

A galvanic cell consists of two beakers. One beaker contains a strip of lead immersed in aqueous sulfuric acid, and the other contains a platinum electrode immersed in aqueous nitric acid. The two solutions are connected by a salt bridge, and the electrodes are connected by a wire. Current begins to flow, and bubbles of a gas appear at the platinum electrode. The spontaneous redox reaction that occurs is described by the following balanced chemical equation: $$3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 8 \mathrm { H } ^ { + } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 4 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$$ Which of the following is the oxidation half-reaction that occurs in the anode?

A) $$\mathrm { Pb } ( \mathrm { s } ) \rightarrow \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$$
B) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } )$
C) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 4 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } ) + 2 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
D) $\mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } ) \rightarrow \mathrm { PbO } ( \mathrm { s } ) + 2 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$
E) $3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 3 \mathrm { e } ^ { - }$

Question 8

An apparatus that is used to generate electricity from a spontaneous redox reaction or, conversely, that uses electricity to drive a nonspontaneous redox reaction is called a(n) _____.

Question 9

Measured redox potentials are independent of the temperature of the system.

Question 10

In any redox reaction, the number of electrons lost by the reductant equals the number of electrons gained by the oxidant.

Question 11

A galvanic cell consists of two beakers. One beaker contains a strip of lead immersed in aqueous sulfuric acid, and the other contains a platinum electrode immersed in aqueous nitric acid. The two solutions are connected by a salt bridge, and the electrodes are connected by a wire. Current begins to flow, and bubbles of a gas appear at the platinum electrode. The spontaneous redox reaction that occurs is described by the following balanced chemical equation: $$3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 8 \mathrm { H } ^ { + } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 4 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$$ Which of the following is the cell diagram for the galvanic cell?

A) $\mathrm{Pt}(\mathrm{s})\left|\mathrm{Pb}^{2+}(\mathrm{aq}) \mathrm{PH}\mathrm{NO}_{3}(\mathrm{aq})\right| \mathrm{NO}(\mathrm{g}) \mid \mathrm{Pb}(\mathrm{s})$
B) $\mathrm { Pb } ( \mathrm { s } ) \left| \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) \mathrm { PNO } ( \mathrm { g } ) \right| \mathrm { Pt } ( \mathrm { s } )$
C) $\mathrm { Pt } ( \mathrm { s } ) \mid \mathrm { Pb } ^ { 2 + } \text { (aq) } \mathrm { PHNO } _ { 3 } ( \mathrm { aq } ) \mid \mathrm { Pb } ( \mathrm { s } )$
D) $\mathrm { Pb } ( \mathrm { s } ) \left| \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) \right| \mathrm { NO } ( \mathrm { g } ) \mathrm { PHNO } _ { 3 } ( \mathrm { aq } ) \mid \mathrm { Pt } ( \mathrm { s } )$
E) $\mathrm { Pb } ( \mathrm { s } ) \left| \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) \mathrm { PHNO } _ { 3 } ( \mathrm { aq } ) \right| \mathrm { NO } ( \mathrm { g } ) \mid \mathrm { Pt } ( \mathrm { s } )$

Question 12

A galvanic cell consists of two beakers. One beaker contains a strip of lead immersed in aqueous sulfuric acid, and the other contains a platinum electrode immersed in aqueous nitric acid. The two solutions are connected by a salt bridge, and the electrodes are connected by a wire. Current begins to flow, and bubbles of a gas appear at the platinum electrode. The spontaneous redox reaction that occurs is described by the following balanced chemical equation: $$3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 8 \mathrm { H } ^ { + } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 4 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$$ Which of the following is the reduction half-reaction that occurs in the cathode?

A) $\mathrm { Pb } ( \mathrm { s } ) \rightarrow \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$
B) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } )$
C) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 4 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } ) + 2 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
D) $\mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } ) \rightarrow \mathrm { PbO } ( \mathrm { s } ) + 2 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$
E) $3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 3 \mathrm { e } ^ { - }$

Question 13

_____ is the study of the relationship between electricity and chemical reactions.

Question 14

A(n) _____ is an ionic substance or solution that allows ions to transfer between the electrode compartments, thereby maintaining the system's electrical neutrality.

Question 15

The standard cell potential is the sum of the tabulated reduction potentials of the two half-reactions.

Question 16

Compare and contrast galvanic cells and electrolytic cells.

Question 17

The _____ of the cell is the difference in electrical potential between two half-reactions.

Question 18

The identity of the salt in a salt bridge is unimportant, as long as the component ions do not react or undergo a redox reaction under the operating conditions of the cell.

Question 19

A(n) _____ is a U-shaped tube inserted into both solutions of a galvanic cell that contains a concentrated liquid or gelled electrolyte and completes the circuit between the anode and the cathode.

Question 20

_____ are the reactions that represent either the oxidation half or the reduction half of a redox reaction.

Question 21

The cell potential values are independent of the stoichiometric coefficients for the half-reaction.

Question 22

If the of the zinc anode is -0.76 V and the of the copper cathode is 0.28 V, what is the of the galvanic cell?

A) 0.480 V
B) 0.760 V
C) -1.04 V
D) -0.480 V
E) 1.04 V

Question 23

The physical states of the reactants and the products in a half-reaction must be identical to those in the overall reaction.

Question 24

The _____ is universally used as a reference electrode and is assigned a standard potential of 0V.

Question 25

Explain the process of measuring standard electrode potentials.

Question 26

Which of the following elements is most likely to be the strongest oxidant?

A) Chlorine
B) Hydrogen peroxide
C) Fluorine
D) Lead
E) Oxygen

Question 27

Explain the use of reference electrodes with examples.

Question 28

Which of the following species can oxidize MnO₂(s) to $\mathrm { MnO } _ { 4 } { } ^ { 2 - }$ under standard conditions?

A) Sn⁴⁺
B) Sn²⁺
C) Al³⁺
D) Li⁺
E) Fe³⁺

Question 29

Which of the following elements is most likely to be the strongest reductant in aqueous solution?

A) Li⁺
B) Al³⁺
C) Zn²⁺
D) Sn⁴⁺
E) Cu²⁺

Question 30

_____ electrodes are designed so that their potential depends on only the concentration of the desired species.

Question 31

All reactants that lie above the standard hydrogen electrode in the table are stronger reductants than H⁺.

Question 32

For the reaction $\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + \mathrm { I } ^ { - } ( \mathrm { aq } ) \rightarrow \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + \mathrm { I } _ { 2 } ( \mathrm { aq } )$ , what is the final balanced equation using half-reactions?

A) $\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + 14 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 6 \mathrm { I } ^ { - } ( \mathrm { aq } ) \rightarrow 2 \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + 7 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } ) + 3 \mathrm { I } _ { 2 } ( \mathrm { aq } )$
B) $\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + 14 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 6 \mathrm { e } ^ { - } ( \mathrm { aq } ) \rightarrow 2 \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + 7 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
C) $\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + 14 \mathrm { H } ^ { + } ( \mathrm { aq } ) \rightarrow 2 \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + 7 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
D) $\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + 2 \mathrm { I } ^ { - } ( \mathrm { aq } ) \rightarrow 2 \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + \mathrm { I } _ { 2 }$
E) $$\mathrm { Cr } _ { 2 } \mathrm { O } _ { 7 } { } ^ { 2 - } ( \mathrm { aq } ) + 2 \mathrm { I } ^ { - } ( \mathrm { aq } ) \rightarrow 2 \mathrm { Cr } ^ { 3 + } ( \mathrm { aq } ) + \mathrm { I } _ { 2 } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$$

Question 33

_____ is defined as the potential of a cell measured under standard conditions.

Question 34

When using a galvanic cell to measure the concentration of a substance, the potential of the _____ is related to the concentration of the substance being measured.

Question 35

The reference electrode should be physically isolated from the solution of interest.

Question 36

The potential of a half-reaction measured against the sunder standard conditions is called the _____ for that half-reaction.

Question 37

If the $E ^ { \circ }$ of the galvanic cell is 0.98 V and the $E ^ { \circ }$ of the copper anode is -0.34 V, what is the $E ^ { \circ }$ of the hydrogen cathode?

A) 1.32 V
B) 0.34 V.
C) -0.64 V
D) 0.64 V
E) 0.34 V

Question 38

If the $E ^ { \circ }$ of the galvanic cell is 1.24 V and the $E ^ { \circ }$ of the cathode is 0.57, what is the $E ^ { \circ }$ of the anode?

A) -0.67 V
B) -0.57 V
C) 1.81 V
D) -1.81 V
E) 0.67 V

Question 39

Which of the following species can reduce PbSO₄ to Pb under standard conditions?

A) Sn²⁺
B) Fe²⁺
C) Cu²⁺
D) Ni²⁺
E) I₂

Question 40

Any species on the left side of a half-reaction will spontaneously oxidize any species on the right side of another half-reaction that lies above it in the table.

Question 41

Derive the relationship between cell potential and the equilibrium constant?

Question 42

If work is being done by a system on its surroundings, it is expressed as a positive number.

Question 43

If the overall cell potential is equal to 1.56 V and two electrons are transferred in the overall reaction, what is the equilibrium constant for the reaction? Given: F = $96,468 \mathrm {~J} / \mathrm { V } \cdot \mathrm { mol } )$ .

A) 1.54×10⁵³
B) 2.78×10⁵²
C) 3.87×10⁵³
D) 5.74×10⁵²
E) 6.98×10⁵³

Question 44

If the equilibrium constant for the reaction is equal to 4.70 ×10⁷⁰ and three electrons are transferred in the overall reaction, what is the overall cell potential? Given: F = $96,468 \mathrm {~J} / \mathrm { V } \cdot \mathrm { mol } )$ .

A) 1.39 V
B) -2.84 V
C) 2.56 V
D) -1.25 V
E) 0.76 V

Question 45

If the reaction 2Ce⁴⁺(aq) + 2Cl^?(aq) ? 2Ce³⁺(aq) + Cl₂(g) is likely to occur spontaneously under normal conditions, what is the $E _ { \text {cell } } ^ { \circ }$ for the reaction?(Use data in the Table 19.2)

A) -0.36V
B) 1.72 V
C) 1.36 V
D) 3.08 V
E) 0.36 V

Question 46

_____ is the fundamental SI unit of electric current.

Question 47

If $E _ { \text {cell } }$ = - 0.76 V, what is the value of the solubility product of the salt?

A) 2.7×10^-12
B) 1.4×10^-13
C) 3.9×10^-13
D) 5.6×10^-12
E) 7.2×10^-13

Question 48

The potentials of two reductive half-reactions can be added to obtain the potential of a third reductive half-reaction.

Question 49

The amount of energy available to do electrical work depends on the total number of electrons that are transferred from the reductant to the oxidant during the course of a reaction.

Question 50

If the cell potential is equal to 0.240 V and the total charge transferred during the reaction is equal to 1.93×10⁵
$\text{C/mol}^{-}$ , what is the maximum work produced by the electrochemical cell?

A) 67.2 kJ/mol
B) -46.3 kJ/mol
C) 27.8 kJ/mol
D) 80.5 kJ/mol
E) -12.9 kJ/mol

Question 51

If the value of Q = 2.5×10³² and three electrons are transferred in the overall reaction, what is the value of $E _ { \text {cell } } ^ { \circ }$ ? Given: E_cell = 0.

A) 0.63 V
B) -0.71 V
C) -0.38 V
D) 0.28 V
E) 0.45 V

Question 52

_____ is the SI unit of measure for the number of electrons that pass a given point in one second.

Question 53

If the free-energy change for the reaction is -87.5 kJ and the cell potential is equal to 0.18 V, how many moles of electrons are used in the reaction? Given: F = $96,468 \mathrm {~J} / (\mathrm { V } \cdot \mathrm { mol } )$ .

A) Two
B) Three
C) Four
D) Five
E) Six

Question 54

If $E _ { \text {cell } } ^ { * }$ = 0.58 V and two electrons are transferred in the overall reaction, what is the value of Q? Given: E_cell = 0.

A) 2.5×10¹⁸
B) 1.9×10¹⁹
C) 5.4×10¹⁸
D) 7.2×10¹⁹
E) 4.2×10¹⁹

Question 55

If $E _ { \text {catrode } } ^ { \circ }$ = 1.60 V and $E _ { \text {arode } } ^ { \circ }$ = 1.25 V, what is the free energy change for the reaction? Given: Total charge transferred during the reaction is equal to 1.93×10⁵ $\text{C/mol}^{-}$ .

A) -74.0 kJ/mol
B) -12.3 kJ/mol
C) -67.6 kJ/mol
D) -28.4 kJ/mol
E) -50.1 kJ/mol

Question 56

Explain the process of determining the oxidants and reductants based on the standard potentials of half-reactions.

Question 57

If the cell potential is equal to 0.760 V and the total charge transferred during the reaction is equal to 2.89×10⁵
$\text{C/mol}^{-}$ , what is the change in free-energy for the reaction?

A) 3.88 *10² kJ/mol
B) -1.32 * 10² kJ/mol
C) -4.29 *10² kJ/mol
D) 6.16 * 10² kJ/mol
E) -2.20 * 10² kJ/mol

Question 58

Multiplying the charge on the electron by Avogadro's number gives us the charge on 1 mol of electrons, which is called the _____.

Question 59

A galvanic cell is constructed with a standard Fe/Fe²⁺ couple in one compartment and a modified hydrogen electrode in the second compartment. The pressure of hydrogen gas is 1.0 atm, but [H⁺] in the second compartment is unknown. What is the pH of the solution in the second compartment if the measured potential in the cell is 0.17 V at $25 ^ { \circ } \mathrm { C }$ ?

A) 3.4
B) 1.9
C) 6.8
D) 8.2
E) 4.7

Question 60

For the reaction $2 \mathrm { Fe } ^ { 3 + } + 2 \mathrm { Cl } ^ { - } \rightarrow 2 \mathrm { Fe } ^ { 2 + } + \mathrm { Cl } _ { 2 }$ , $E _ { \text {cell } } ^ { \circ }$ = -0.59 V. What is the E for the reaction? Given: [Fe³⁺] = 0.017 M, [Fe²⁺] = 0.50 M, [Cl^-] = 0.035 M, $\mathrm { P } _ { \mathrm { Cl } _ { 2 } }$ = 1 atm, and T = 298 K.

A) 0.34 V
B) -0.76 V
C) -0.28 V
D) 0.57 V
E) 0.82 V

Question 61

Explain the Leclanche dry cell and state its half-reactions.

Question 62

In a lithium-iodine battery, the anode is a solid complex of I₂ and the cathode is lithium metal.

Question 63

Explain fuel cells with an example. State its half-reactions.

Question 64

The alkaline battery has a longer shelf life and more constant output voltage than the Leclanche dry cell.

Question 65

Which of the following is the reaction that occurs at the cathode of a button battery?

A) 2MnO₂(s) + 2NH₄⁺(aq) + 2e^? ? Mn₂O₃(s) + 2NH₃(aq) + H₂O(l)
B) I₂(s) + 2e^? ? 2I^?(LiI)
C) 2NiO(OH)(s) + 2H₂O(l) + 2e^? ? 2Ni(OH)₂(s) + 2OH^?(aq)
D) PbO₂(s)+ $\mathrm { HSO } _ { 4 } ^ { - }$ (aq)+3H⁺(aq)+2e ^- ?PbSO₄(s)+2H₂O(l)
E) HgO(s) + H₂O(l) + 2e^? ? Hg(l) + 2OH^?(aq)

Question 66

The _____ is a battery that consists of a Leclanché cell adapted to operate under alkaline conditions.

Question 67

The anode of each cell in a lead storage battery is a grid of spongy lead metal, and the cathode is a similar grid containing powdered lead dioxide.

Question 68

The NiMH battery is more advantageous than the NiCad battery in that _____.

A) it has a lower self-discharge rate
B) it is insensitive to recharging memory
C) it has an unlimited service life
D) it requires lower maintenance
E) it is less expensive

Question 69

An electrochemical cell in which the anode and cathode compartments are identical except for the concentration of a reactant is called a(n) _____.

Question 70

During cathodic protection, a more reactive metal such as zinc becomes the cathode and iron becomes the anode.

Question 71

Which of the following is the half-reaction that occurs at the cathode of the Leclanche dry cell?

A) PbO₂ (s)+ $\mathrm { HSO } _ { 4 } ^ { - }$ (aq)+3H⁺(aq)+2e^- ?PbSO₄(s)+2H₂O(l)
B) 2NiO(OH)(s) + 2H₂O(l) + 2e^? ? 2Ni(OH)₂(s) + 2OH^?(aq)
C) HgO(s) + H₂O(l) + 2e^? ? Hg(l) + 2OH^?(aq)
D) 2MnO₂(s) + 2NH₄⁺(aq) + 2e^? ? Mn₂O₃(s) + 2NH₃(aq) + H₂O(l)
E) Ag₂O(s) + H₂O(l) + 2e^? ? 2Ag(s) + 2OH^?(aq)

Question 72

Which battery is used to provide the starting power in virtually every automobile and marine engine?

A) Lead-acid battery
B) Nickel-cadmium battery
C) Lithium-iodine battery
D) Leclanche dry cell
E) Alkaline battery

Question 73

Which of the following reactions occurs at the cathode during the formation of rust?

A) $\mathrm { O } _ { 2 } ( g ) + 4 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 4 \mathrm { e } ^ { - } \rightarrow 2 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
B) $\mathrm { Fe } ( \mathrm { s } ) \rightarrow \mathrm { Fe } ^ { 2 + } ( \mathrm { aq } ) + 2 e ^ { - }$
C) Zn(s) ? Zn²⁺(aq) + 2e^?
D) 2H₂(g) ? 4H⁺ + 4e^?
E) $\mathrm { PbO } _ { 2 } ( \mathrm {~s} ) + \mathrm { HSO } _ { 4 } ^ { - } ( \mathrm { aq } ) + 3 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - } \rightarrow \mathrm { PbSO } _ { 4 } ( \mathrm {~s} ) + 2 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$

Question 74

Which of the following is the overall reaction of the Leclanche dry cell?

A) Zn(s) + 2HgO(s)→Hg(l)+ZnO(s)
B) 2MnO₂(s) + 2NH₄Cl(aq) + Zn(s) → Mn₂O₃(s) + Zn(NH₃)₂Cl₂(s) + H₂O(l)
C) Zn(s) + 2Ag2O(s)→2Ag(s)+ZnO(s)
D) Cd(s)+2NiO(OH)(s)+2H2O(l)→Cd(OH)2(s)+2Ni(OH)2(s)
E) Pb(s) + PbO₂(s)+2HSO4−(aq)+2H+(aq)→2PbSO4(s)+2H2O(l)

Question 75

A battery is a galvanic cell that requires a constant external supply of one or more reactants to generate electricity.

Question 76

Which of the following metals is most likely to protect iron from corrosion?

A) Sn²⁺
B) Ni²⁺
C) Cr³⁺
D) Zn²⁺
E) Cu²⁺

Question 77

Which of the following is the overall reaction for the formation of rust?

A) 2MnO₂(s) + 2NH₄Cl(aq) + Zn(s) → Mn₂O₃(s) + Zn(NH₃)₂Cl₂(s) + H₂O(l)
B) Cd(s)+2NiO(OH)(s)+2H₂O(l)→Cd(OH)₂(s)+2Ni(OH)₂(s)
C) 4Fe²⁺(aq) + O₂(g) + (2 + 4x)H₂O → 2Fe₂O₃·xH₂O + 4H⁺(aq)
D) Zn(s) + 2MnO₂(s)→ZnO(s)+Mn₂O₃(s)
E) 2Zn(s) + O₂(g)+4H⁺(aq)→2Zn²⁺(aq)+2H₂O(l)

Question 78

Lithium-iodine batteries are disposable and cannot be recharged once they are discharged.

Question 79

The _____ is a battery consisting of an electrolyte that is an acidic water-based paste containing MnO₂, NH₄Cl, ZnCl₂, graphite, and starch.

Question 80

Which of the following is the overall reaction of a fuel cell?

A) Pb(s) + PbO₂(s)+2HSO₄^- (aq)+2H⁺(aq)→2PbSO₄(s)+2H₂O(l)
B) Cd(s)+2NiO(OH)(s)+2H₂O(l)→Cd(OH)₂(s)+2Ni(OH)₂(s)
C) 2H₂(g)+O₂(g)→2H₂O(g)
D) 2Li(s) + I₂(s)→2LiI(s)
E) 2MnO₂(s) + 2NH₄Cl(aq) + Zn(s) → Mn₂O₃(s) + Zn(NH₃)₂Cl₂(s) + H₂O(l)