Chapter 18

Question 1

Non-Standard State Solutions

Answer 1

∆G = ∆G^o + RTln(Q)

Question 2

Electrochemical Cell

Answer 2

A device in which a chemical reaction either produces or is carried out by an electrical current.

Question 3

Voltaic (Galvanic Cell)

Answer 3

An electrochemical cell that produces electrical current from a spontaneous chemical reaction.

Question 4

Electrolytic Cell

Answer 4

Consumes electrical current to drive a nonspontaneous chemical reaction.

Question 5

Half-Cell

Answer 5

One half of an electrochemical cell where either oxidation or reduction occurs.

Question 6

Electrodes

Answer 6

Conductive surfaces through which electrons can enter or leave the half-cells.

Question 7

Amperes (A)

Answer 7

• Measure of electrical current.
• 1 A = 1 C/s

Question 8

Potential Difference

Answer 8

A measure of the difference in potential energy per unit of charge (J/C)

1 J/C = V

Question 9

Electromotive Force (emf)

Answer 9

The force that results in the motion of electrons due to a difference in potential.

Question 10

Cell Potential/Cell emf (E_cell)

Answer 10

The potential difference between the two electrodes (in a voltaic cell)

• Depends on:

1. Depends on tendencies of reactants to undergo oxidation and reduction.
2. Depends on the concentrations of the reactants and products in the cell.
3. Temperature.

Question 11

Standard Cell Potential / Standard emf (E^o_cell)

Answer 11

• I M concentration for reactants in solution
• 1 atm for pressure for gaseous reactants
• Temperature: 25^oC

Question 12

Anode

Answer 12

• Electrode where oxidation occurs.
• More negatively charged electrode.

Question 13

Cathode

Answer 13

• Electrode where reduction occurs.
• More positivetly charged electrode.

Question 14

Salt Bridge

Answer 14

• Inverted, U-shaped tube that contains a strong electrolyte such as KNO₃ and connects the two half-cells.
• Allows a flow of ions that neutralize the charge buildup in the solution.
• The (-) ions within the salt bridge flow to neutralize the accumulation of postive charge at the anode, and the (+) ions flow to neutralize the accumulation of negative charge at the cathode (The salt bridge completes the circuit).

Question 15

Cell Diagram / Line Notation

Answer 15

• Oxidation half-reaction written on the left.
• Reduction half-reaction written on the right.
• Substances in different phases are separated by a single vertical line.
• Reactant and products of one or both of the half-reactions may be in the same phase, where we then separate them from each other with a comma.
• A double vertical line (salt bridge) separates the two half-reactions.

Example:

Zn(s)|Zn²⁺(aq)||Cu²⁺(aq)|Cu(s)

Question 16

Standard Electrode Potential

Answer 16

The overall standard cell potential (E^o_cell) is the difference between the two stard electrode potentials.

Question 17

Standard Hydrogen Electrode (SHE)

Answer 17

Half-cell consisting of an intert platinum electrode immersed in 1 M HCl with hydrogen gas at 1 atm bubbling through the solution; used as the standard of a cell potential of zero.

2H⁺(aq) + 2e^- → H₂(g), E^o_cathode = 0.00 V

Question 18

E^o_cell (formula)

Answer 18

E^o_cell = E^o_final - E^o_initial

= E^o_cathode - E^o_anode

Question 19

Standard Electrode Potentials (Summary)

Answer 19

• The electrode potential of the SHE is exactly zero.
• The electrode in any half cell with a greater tentency to undergo reduction is positively charged relative to the SHE and therefore has a positive E^o.
• The electrode in any half-cell with a lesser tendency to undergo reduction (or greater
  tendency to undergo oxidation) is negatively charged relative to the SHE and therefore has a negative E^o.
• The cell potential of any electrochemical cell (Eocell) is the difference between the
  electrode potentials of the cathode and the anode (E^o_cell = E^o_cathode - E^o_anode).
• E^o_cell is positive for spontaneous reactions and negative for nonspontaneous reactions.

E°_cell is positive for spontaneous reactions and negative for nonspontaneous reactions.

Question 20

Prediction of Spontaneous Direction for Redox Reactions

Answer 20

• The half-reaction with the more positive electrode potential attracts electrons more
  strongly and will undergo reduction.
• The half-reaction with the more negative electrode potential repels electrons more strongly and will undergo oxidation.

Question 21

Spontaneous Reactions

Answer 21

• Proceeds forward when all reactants and products are in their standard states.
• ΔG^o is negative (<0)
• E^o_cell is positive (>0)
• K > 1

Question 22

Nonspontaneous Reaction

Answer 22

• Proceeds in the reverse direction when all reactants and products are in their standard states.
• ΔG^o is positive (>0)
• E^o_cell is negative (<0)
• K < 1

Question 23

Standard Change in

Free Energy (ΔG^o)

Answer 23

ΔG^o = -nFE^o_cell

• is the standard change in free enthalpy ΔG^o for an electrochemial reaction.
• n is the number of moles of electrons transferred in the balanced equation.
• F = Faraday’s Constant.
• E^o_cell is the standard cell potential.
  *

Question 24

Faraday’s Constant (F)

Answer 24

Represents the charge in coulombs of 1 mol of electrons.

F = 96,485 C / 1 mol e^-

Question 25

Relationship between E^o_cell and K

Answer 25

E^o_cell = (0.0592 V / n)\*log(K)

Question 26

Nernst Equation

Answer 26

E_cell = E^o_cell - (0.0592 V / n)\*log(Q)

Question 27

Dry-Cell Batteries

Answer 27

A battery that does not contain a large amount of liquid water, often using the oxidation of zinc and the reduction of MnO₂ to provide the electrical current.

Question 28

Lead-Acid Storage Batteries

Answer 28

A battery that uses the oxidation of lead and the reduction of lead(IV) oxide in sulfuric acid to provide electrical current.

Question 29

Nickle-Cadmium (NiCad) Battery

Answer 29

Consists of an anode composed of solid cadmium and a cathode composed of Ni(OH) (s) in a KOH solution.

Question 30

Nickel-Metal Hydride (NiMH) Battery

Answer 30

A battery that uses the same cathode reaction as the NiCad battery but a different anode reaction, the oxidation of hydrogens in a metal alloy.

Question 31

Fuel Cells

Answer 31

* The reactants (the fuel provided by an external source) constantly flow through the battery, generating electrical current as they undergo a redox reaction.

Question 32

Electrolysis

Answer 32

The process by which electrical current is used to drive an otherwise nonspontaneous redox reaction.

Question 33

Characteristics of Electrochemical Cell Types

Answer 33

* In all electrochemical cells: * Oxidation occurs at the anode. * Reduction occurs at the cathode. * In voltaic cells: * The anode is the source of electrons and has a negative charge (anode -) * The cathode draws electrons and has a positive charge (cathod +) * In electrolytic cells: * Electrons are drawn away from the anode, which must be connected to the positive terminal of the external power source (anode +) * Electrons are forced to the cathode, which must be connected to the negative terminal of the power source (cathode -)

Question 34

Electrolysis of Pure Molten Salts

Answer 34

The anion is oxidized and the cation is reduced.

Question 35

Electrolysis of mixtures of cations or anions

Answer 35

* The cation that is most easily reduced (the one with the more positive elctrode potential) is reduced first. * The anion that is most easily oxidized (the one with the more negative electrode potential) is oxidized first.

Question 36

Electrolysis of aqueous solutions

Answer 36

The cations of active metals - those that are not easily reduced (Li⁺, K⁺, Na⁺, MG²⁺, Ca²⁺, Al³⁺) - cannot be reduced from aqueous solutions by electrolysis because water is reduced at a lower voltage.

Question 37

Corrosion

Answer 37

Oxidation of metals that occurs when they are exposed to oxidizing agents in the environment.

Question 38

Formation of Rust

Answer 38

* *Moisture must be present for rusting to occur as water is a reactant.* * *Additional electrolytes promote rusting* as their presence on the surface of iron promotes rusting because it enhances current flow. * *The presence of acids promotes rusting*. Since H⁺ ions are involved in the reduction of oxygen, lower pH enhances the cathodic reaction and leads to faster rusting.

Question 39

Oxidation

Answer 39

The loss of one or more electrons. The gaining of oxygen or the loss of hydrogen.

Question 40

Reduction

Answer 40

The gaining of one or more electrons. The gaining of hydrogen or loss of oxygen.

Question 41

Balancing Aqueous Redox Rxn's in acidic solution

Answer 41

1. Separate the overall reactions into two half-reactions: oxidation and reduction. 2. Balance each half-reaction in the following order: * Balance all other elements other than H and O * Balance O by adding H₂O * Balance H by adding H⁺ 3. Balance each half-reaction with respect to charge by adding electrons. 4. Make the number of electrons equal by multiplying one/both half-reactions. 5. Add the two half-reactions together, canceling out electrons and other species as necessary.

Question 42

Balancing Redox Reactions Occurring in Basic Solution

Answer 42

1. Separate the overall reactions into two half-reactions: oxidation and reduction. 2. Balance each half-reaction with respect to mass: * Balance all elements other than H and O * Balance H by adding H⁺ * Neutralize H⁺ by adding enough OH^- to neutralize each H⁺. Add the same number of OH^- ions to each side of the equation. 3. Balance each half-reaction with respect to charge by adding electrons. 4. Make the number of electrons in both half-reactions equal. 5. Add the two half-reactions together, canceling out electrons and other species as necessary.