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Gen Chem > Ch. 12: Electrochemistry (Complete) > Flashcards

Ch. 12: Electrochemistry (Complete) Flashcards

1
Q

defn: electrochemical cells + 3 fundamental types

A

contained systems in which oxidation-reduction reactions occur

galvanic cells (voltaic cells)
electrolytic cells
concentration cells

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2
Q

what are 5 commonalities between the 3 types of electrochemical cells?

A
  1. all contain electrodes where oxidation and reduction take place
  2. all have a reduction reaction at cathode
  3. an oxidation reaction at anode
  4. current flowing from cathode to anode
  5. electron flow from anode to cathode
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3
Q

what are two differences between the 3 types of electrochemical cells?

A

galvanic and concentration: house spontaneous reaction, connected by a conductive material/housed in different compartments

electrolytic cells: contain nonspontaneous reactions, housed in the same compartment

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4
Q

defn + mnemonic: anode + cathode

A

ANODE = electron where oxidation occurs

CATHODE = electron where reduction occurs

AN OX, RED CAT

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5
Q

defn: electromotive force (emf)

what does it indicate when emf is positive? negative?

A

corresponds to the voltage or electrical potential difference of the cell

emf positive = cell can release energy (delta G < 0) = spontaneous

emf negative = cell must absorb energy (delta G > 0) = nonspontaneous

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6
Q

what is the direction of electrons and of current for all electrochemical cells?

A
  • movement of electrons: anode to cathode
  • current (I) runs cathode to anode
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7
Q

what is the relationship between the flow of electrons and the current?

A

they are always of equal magnitude but in opposite directions

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8
Q

are batteries influenced by temperature? how?

A

yes!

most galvanic cells fail in cold weather

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9
Q

what is the relationship between the free energy change and the emf?

A

they always have opposite signs

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10
Q

char (4): galvanic/voltaic cells

A
  1. all nonrechargeable household batteries you own
  2. spontaneous reactions
  3. delta G < 0 (the reactions free energy is decreasing) as the cell releases energy to the environment
  4. Ecell (emf) = positive
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11
Q

what is the basic setup of a galvanic/voltaic cell? (5)

A
  1. two electrodes of distinct chemical identity are placed in separate compartments (called half-cells)
  2. the two electrodes are connected to each other by a conductive material (like copper wire)
  3. along the wire, there may be other various circuit components (i.e. resistors, capacitors, etc.)
  4. surrounding each electrode is an aqueous electrolyte solution composed of cations and anions
  5. connecting the two solutions is a salt bridge (consists of an inert salt)
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12
Q

can the cations in the two half-cell solutions be of the same element as the respective metal electrode in a galvanic/voltaic cell?

A

yes!

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13
Q

what is the process that occurs when the electrodes in a galvanic/voltaic cell are connected to each other by a conductive material? (4)

A
  1. charge will begin to flow as the result of an redox reaction that is taking place between the two-half cells
  2. the redox reaction is spontaneous, so the change in Gibbs free energy is negative
  3. as the spontaneous reaction proceeds towards equilibrium, the movement of electrons results in a conversion of electrical potential energy into kinetic energy
  4. by separating the reduction and oxidation half-reactions into two compartments, we are able to harness this energy and use it to do work by connecting various electrical devices into the circuit between the two electrodes
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14
Q

Daniell cell: defn

set up

what are the anode and cathode?

A

a specific galvanic/voltaic cell

a zinc electrode is placed in an aqueous ZnSO4 solution
a copper electrode is placed in an aqueous CuSO4 solution

anode: zinc bar (Zn (s) is oxidized to Zn2+ (aq))
cathode: copper bar (Cu 2+ (aq) is reduced to Cu (s))

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15
Q

what are the half-cell reactions, net reaction, and Ered/Ecell for each reaction?

A

Zinc: Ered = -0.762 V (anode)
Copper: Ered = +0.340 V (cathode)

Net: Ecell = +1.102 V

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16
Q

what direction do electrons flow in the Daniell cell?

A

from the zinc anode through the wire to the copper cathode

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17
Q

what flows external from the salt bridge into the Daniell cell? (2)

A

anions (Cl-) flow externally from the salt bridge into the ZnSO4
cations (K+) flow externally from the salt bridge into the CuSO4

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18
Q

what would happen if the two half-cells were not seperated?

A

the Cu2+ ions would react directly with the zinc bar, and no useful electrical work would be done

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19
Q

what would happen if only a wire were provided for this electron flow?

A

the reaction would soon stop because an excess positive charge would build up on the anode and an excess negative charge would build up on the cathode

the excessive charge accumulation would provide a countervoltage large enough to prevent the redox reaction from taking place and the current would cease

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20
Q

func: salt bridge

A

to exchange of anions and cations to balance/dissipate newly generated charges

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21
Q

what does the salt bridge contain?

A

an inert electrolyte (KCl or NH4NO3 usually) which contains ions that will not react with the electrodes or with the ions in solution

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22
Q

explain how the salt bridge works in the context of the Daniell cell

A

the anions from the salt bridge (Cl-) diffuse into the solution on the anode side (ZnSO4) to balance out the charge of the newly created Zn2+ ions

the cations of the salt bridge (K+) flow into the solution on the cathode side (CuSO4) to balance out the charge of the sulfate ions left in solution when the Cu2+ ions are reduced to Cu and precipitate onto the electrode

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23
Q

defn: plating or galvanization

A

the precipitation process of the reduced Cu (in the case of the Daniell cell) onto the cathode

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24
Q

what accounts for the relatively short lifespan of the Daniell cell?

A

the flow of anions and cations from the salt bridge into the half-cells and the finite quantity of Cu2+ in the solution

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25
mnemonic: flow of electrons in all electrochemical cells
A --> C alphabetical electrons flow from Anode to Cathode
26
defn + example for Daniell cell: cell diagram
shorthand notation representing the reactions in an electrochemical cell Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
27
what are the 3 rules for constructing a cell diagram?
1. reactants and products are always listed from L to R in this form: anode | anode solution (concentration) || cathode solution (concentration) | cathode 2. a single vertical line indicates a phase boundary 3. a double vertical line indicates the presence of a salt bridge or some other type of barrier
28
char: electrolytic cell (3)
1. house nonspontaneous reactions that require the input of energy to proceed 2. delta G > 0 3. Emf negative
29
defn: electrolysis
oxidation-reduction reaction driven by an external voltage source, and in which chemical compounds are decomposed
30
what materials are electrodes made of in electrolytic cells?
any material so long as it can resist the high temperatures and the process' corrosion
31
explain the process of what happens in an electrolytic cell (2) in this example: molten NaCl is decomposed into Cl2 (g) and Na (l)
1. the external voltage source (a battery) supplies energy sufficient to drive the redox reaction in the direction that is thermodynamically unfavorable (nonspontaneous) 2. Na+ ions migrate toward the cathode, where they are reduced to Na (l). Cl- ions migrate toward the anode, where they are oxidized to Cl2
32
why do the half reactions in an electrolytic cell not need to be separated?
because the desired reaction is nonspontaneous
33
what do Faraday's laws state?
the liberation of gas and deposition of elements on electrodes is directly proportional to the number of electrons being transferred during the redox reaction
34
what is the equation and interpretation of the equation that corresponds to this concept?
M^(N+) + ne- --> M(s) one mole of metal M (s) will logically be produced if n moles of electrons are supplied to one mole of M^(n+)
35
what charge does one electron carry?
1.6 * 10^-19 coulombs (C)
36
defn + values: Faraday constant / faraday (F)
one faraday (F) is equivalent to the amount of charge contained in one mole of electrons (1 F = 96,485 C) or one equivalent Faraday constant = 96485 C/mol e- (round to 10^5 C/mole e-)
37
eqn + mnemonic: electrodeposition equation
calculating MOLES OF Metal, IT is Not Fun where mol M = amount of metal ion being deposited at a specific electrode I = current t = time n = number of electron equivalents for a specific metal ion F = Faraday constant
38
func: electrodeposition equation
- helps determine the number of moles of element being deposited on a plate - can be used to determine the amount of gas liberated during electrolysis
39
are delta G and emf the same sign or opposite signs in electrochemical cells?
opposite
40
char (2): concentration cell
1. a special type of galvanic cell 2. contains 2 half-cells connected by a conductive material (like all galvanic cells), which allows a spontaneous redox reaction proceeds (generates a current and delivers energy)
41
what is the distinguishing characteristic of a concentration cell?
the electrodes are chemically identical
42
what implication does this have on how the cell functions?
both electrodes have the same reduction potential
43
how is current generated in a concentration cell?
as a function of a concentration gradient established between the two solutions surrounding the electrodes
44
what drives the movement of electrons?
the concentration gradient results in a potential difference between the two compartments, drives the movement of electrons in the direction that results in equilibration of the ion gradient
45
when does the current stop in a concentration cell? what is the implication of this?
when the concentrations of ionic species in the half-cells are equal this implies that the voltage or electromotive force of a concentration cell is 0 when the concentrations are equal
46
how can the voltage be calculated in concentration cells?
using the Nernst equations as a function of concentrations
47
defn: rechargeable cell / rechargeable battery
one that can function as both a galvanic and electrolytic cell
48
defn: lead-acid battery (lead storage battery)
a specific type of rechargeable battery
49
since a lead-acid battery is a voltaic cell, what composes the battery when it is fully charged? fully discharged?
when fully charged: it consists of two half-cells - a Pb anode and a porous PbO2 cathode, connected by a conductive material (concentrated 4 M H2SO4) when fully discharged: it consists of two PbSO4 electroplated lead electrodes with a dilute concentration of H2SO4
50
what are the half-reactions of lead-acid batteries? are their Ered's positive or negative?
oxidation half-reaction at lead (negative) anode Pb (s) + HSO4- (aq) --> PbSO4 (s) + H+ (aq) + 2 e- (negative Ered) reduction half-reaction at lead IV oxide (positive) cathode PbO2 (s) + SO42- (aq) + 4H+ )+ 2e- --> PbSO4 (s) + 2 H2O (positive Ered)
51
what happens when a lead-acid battery is discharging? (2)
part of a voltaic circuit both half-reactions cause the electrodes to plate with lead sulfate and dilute the acid electrolyte
52
what happens when a lead-acid battery is charging? (2)
part of an electrolytic circuit external source reverses the electroplating process and concentrates the acid solution
53
defn: energy density (2)
energy to weight ratio a measure of a battery's ability to produce power as a function of its weight
54
do lead-acid batteries have a high or low energy density? what implication does this have?
low this means that lead-acid batteries require a heavier amount of battery material to produce a certain output as compared to other batteries
55
defn + char: nickel-cadmium batteries (3)
another type of rechargeable cell 1. consist of two half-cells made of solid cadmium (anode) and nickel 3 oxide-hydroxide (cathode) connected by a conductive material (usually potassium hydroxide KOH) 2. higher energy density than lead-acid batteries 3. electrochem of this type of battery provides a high surge current
56
are the AA and AAA batteries we typically use lead-acid or nickel-cadmium?
nickel-cadmium
57
what are the half-reactions of a nickel-cadmium battery? what do both half-reactions do?
oxidation half-reaction at cadmium (negative) anode Cd (s) + 2 OH- (aq) --> Cd(OH)2 (s) + 2e- (E negative) reduction half-reaction at nickel oxide-hydroxide (positive) cathode 2 NiO(OH) (s) + 2H2O + 2e- --> 2 Ni(OH)2 (s) + 2OH- (E positive) both half-reactions cause the electrodes to plate with their respective products
58
why are some Ni-Cd designs vented?
charging reverses the electrolytic cell potentials, so sometimes they are vented to allow for the release of built-up hydrogen and gas during electrolysis
59
defn: surge current
periods of large current (amperage) early in the discharge cycle
60
char (5): nick-metal hydride (NiMH) batteries
1. largely have replaced Ni-Cd batteries because more efficient 2. have more energy density 3. more cost effective 4. less toxic 5. a metal hydride is used in lieu of a pure metal anode
61
describe electrode charge designations of galvanic and electrolytic cells
galvanic cell: anode is negative, cathode is positive electrolytic cell: anode is positive, cathode is negative (because an external source is used to reverse the charge) in both: reduction at cathode, oxidation at anode, cations attracted to the cathode, anions are attracted to the anode, electros always flow through the wire from the anode to the cathode and current flow from cathode to anode
62
2 mnemonics for cathode, anode, cations, anions
RED CAT AN OX = reduction at the cathode, oxidation at the anode ANions are attracted to the ANode CATions are attracted to the CAThode
63
defn: isoelectric focusing
a technique used to separate amino acids or polypeptides based on their isoelectric points (pI)
64
defn + func: standard hydrogen electrode (SHE)
given a potential of 0 V by convention used as a reference point for defining and measuring reduction potentials
65
defn + func: reduction potential
the intrinsic tendency of a species to gain electrons and to be reduced allows us to determine which species in a reaction will be oxidized or reduced
66
defn: standard reduction potential
a reduction potential that is measured under standard conditions (25 deg C, 1 atm pressure, 1 M concentrations)
67
what can the relative reactivities of different half-cells be compared for?
to predict the direction of electron flow
68
what does a more positive standard reduction potential indicate? a less positive one?
more positive: a greater relative tendency for reduction to occur less positive: a greater relative tendency for oxidation to occur
69
explain the spontaneity and cathode/anode specifications in terms of reduction potentials for galvanic cells (3)
1. the electrode with the more positive reduction potential is the cathode 2. the electrode with the less positive reduction potential is the anode 3. because the species with a stronger tendency to gain electrons (that wants to gain electrons) is actually doing so, the reaction is spontaneous and deltaG is negative
70
explain the spontaneity and cathode/anode specifications in terms of reduction potentials for electrolytic cells (3)
1. the electrode with the more positive reduction potential is forced by the external voltage source to be oxidized and is, thus, the anode 2. the electrode with the less positive reduction potential is forced to be reduced and is therefore, the cathode 3. because the movement of electrons is in the direction against the tendency or desires of the respective electrochemical species, the reaction is nonspontaneous and deltaG is positive
71
how do you obtain the oxidation potential of a half-reaction?
both the reduction half-reaction and the sign of the reduction potential are reversed reduction and oxidation potentials have equal magnitudes but opposite signs
72
defn: standard electromotive force (emf)
the difference in potential (voltage) between two half-cells under standard conditions
73
eqn: standard emf
determined by calculating the difference in reduction potentials between the two half-cells
74
when subtracting standard potentials, why shouldn't you multiply them by the number of moles oxidized or reduced?
because the potential of each electrode does not depend on the size of the electrode (the amount of material), but rather the identity of the material
75
should you multiply the reduction potential if you need to multiply each half-reaction by a common denominator to cancel out electrons when coming up with a net ionic equation? why or why not?
no, because that would indicate a change in chemical identity of the electrode, which is not occurring
76
what is the equation that relates the change in gibbs free energy to emf?
where delta Gdeg = the standard change in free energy n = the number of moles of electrons exchanged F = the Faraday constant Edegcell = the standard emf of the cell
77
eqn + func: Nernst equation
how we calculate emf when conditions deviate from standard conditions where Ecell = emf of the cell under nonstandard conditions Edegcell = emf of the cell under standard conditions R = ideal gas constant T = temp in kelvin n = # of moles of electrons F = Faraday constant Q = reaction quotient for the reaction at a given point in time
78
what is the simplified version of the Nernst equation (which should be used on test day), assuming T = 298K?
79
func: voltmeter
measures the emf of a cell
80
defn: potentiometer
a kind of voltmeter that draws no current and gives a more accurate reading of the difference in potential between two electrodes
81
what does a negative Edegcell correlate with in terms of spontaneity of a reaction? what about a positive Edegcell?
negative: nonspontaneous positive: spontaneous
82
eqn: change in Gibbs free energy of an electrochemical cell with varying concentrations

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