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Chem 2 Heider > Ch 18 - Electrochemistry > Flashcards

Ch 18 - Electrochemistry Flashcards

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1
Q

the most common type of fuel cell is the

A

hydrogen-oxygen fuel cell

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2
Q

fuel cells are based on

A

oxidation-reduction reactions

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3
Q

being reduced means

A

you gain more electrons or become more negative

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4
Q

oxidation is the

A

loss of electrons

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5
Q

reduction is the

A

gain of electrons

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6
Q

oxidation corresponds to an increased

A

oxidation state

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7
Q

reduction corresponds to a decrease in

A

oxidation state

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8
Q

both the ______ and ______ must be balanced in an ox-redox reaction

A

mass and the charge

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9
Q

half-reaction method of balancing

A
  • break down the overall equation into two half reactions(oxidation and reduction)
    - balance the two equations individually and then add them together
    - slightly different for acids in basic solution
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10
Q

electrical current

A

the flow of electric charge

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11
Q

electrochemical cell

A

the generation of electricity though a redox reaction is carried out in this device

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12
Q

voltaic(galvanic) cell

A

an electrochemical cell that produces electrical current from spontaneous chemical reactions

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13
Q

electrolytic cell

A

consumes electrical current to drive a nonspontaneous chemical reaction

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14
Q

half-cell

A

splitting a voltaic cell into two parts connected by a wire to control the flow of electricity

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15
Q

electrodes

A

conductive surfaces though which electrons can enter or leave the half cells

	- each metal strip reaches equilibrium with its ions in solution
	- Zn(s)  Zn^2+(aq) + 2e^-
	- Cu(s)  Cu^2+(aq) + 2e^-
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16
Q

the two metal electrodes come to equilibrium at different points which it what drives

A

the electrons in one direction creating electricity

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17
Q

electrical current is measured in

A

amperes(A)

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18
Q

one ampere represents the flow of one coulomb(a measure of electrical charge) per second

A

1A = 1C/s

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19
Q

electron current is driven by the difference is potential energy caused by an electric field resulting from

A

the charge difference of two electrodes

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20
Q

potential difference

A

a measure of the difference in potential energy(usually joules) per unit of a charge(coulombs)
- SI unit of potential difference is the volt(V)

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21
Q

volt(V)

A

one joule per coulomb

	- the SI unit of potential difference
	- 1V = 1J/C
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22
Q

the potential difference of one volt indicates

A

that a charge of one coulomb experiences an energy difference of one joule between the two electrodes

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23
Q

a large potential difference =

A

large difference in charges of electrodes = strong tendency for electron flow

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24
Q

electromotive force(emf)

A

another name for potential difference since it results in a force for the motion of electrons

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25
in a voltaic cell the potential difference between the two electrodes is
the cell potential(Ecell) or cell emf
26
cell potential depends on
the relative tendencies of the reactants to undergo oxidation and reduction - stronger the tendencies the larger the difference and therefore a higher potential
27
cell potential also depends on
the concentrations of the reactants and products and temperature
28
standard cell potential(Edegree symbol cell) or standard emf
the cell potential at standard conditions(1 atm for gaseous reactants and 1M for solutions)
29
cell potential is a measure of
the overall tendency of the redox reaction to occur spontaneously
30
a negative cell potential indicates
that the forward reaction is not spontaneous
31
anode
in all electrochemical cells the electrode where oxidation occurs
32
cathode
in all electrochemical cells the electrode where reduction occurs
33
in a voltaic cell the anode is
the more negatively charged electrode labeled with the negative(-) sign
34
in a voltaic cell the cathode is
the more positively charged electrode and labeled with the positive(+) sign
35
electrons flow from the anode to cathode through wires
from negative to positive
36
as electrons flow out of the anode positive charge builds up in the solution
form in the oxidation half cell
37
as electrons flow into the cathode positive ions are reduced at the reduction half cell
negative charge builds up in the solution
38
salt bridge
an inverted U shaped tube that contains a strong electrolyte such as KNO3 and connects the two half cells
39
the negative ions within the salt bridge flow to neutralize the accumulation of positive charge at the anode and the positive ions flow to neutralize the accumulation of negative charge at the cathode
- the salt bridge completes the circuit allowing electrical current to flow
40
standard cell potential(Edegree sign cell for an electrochemical cell depends on
the specific half reactions occurs and the potential energy difference(per unit charge) of the two electrodes
41
standard electrode potential
each electrode in each half cell has its own individual potential
42
the overall cell potential(Edegree Cell) is
the difference between the two standard electrode potentials
43
while cells are connected electrons flow from the electrode with
more negative charge(greater potential energy) to the electrode with more positive charge(less potential energy)
44
standard hydrogen electrode(SHE)
normally chosen to have a potential of zero - consist of an inert platinum electrode immersed in 1M HCl with hydrogen gas at 1 atm bubbling through the solution - Edegree cell = Edegree final – Edegree initial - Edegree cathode – Edegree anode
45
Edegree cell =
Edegree final – Edegree initial | - Edegree cathode – Edegree anode
46
the electrode potential of the standard hydrogen electrode(SHE) =
0
47
the electrode in any half-cell with a greater tendency to undergo reduction is
positively charged relative to the SHE and therefore has a positive Edegree cell
48
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 Edegree cell
49
the cell potential of any electrochemical cell(Edegree cell) is the difference between
the electrode potentials of the cathode and anode(Edegree cell = Edegree cathode – Edegree anode)
50
Edegree cell is positive for _______ reactions and negative for ______ reactions
spontaneous nonspontaneous
51
NIO
more negative is oxidation
52
PIR
more positive is reduction
53
the half reaction with the more positive electrode potential
attracts electrons more strongly and will undergo reduction
54
the half reaction with the more negative electrode potential
repels electrons more strongly and will undergo oxidation
55
any reaction is table 18.1 is spontaneous when paired with
the reverse of any of the reactions listed below it on the table
56
For a spontaneous redox reaction(in standard states the forward direction)
- deltaGdegree 0 | - K > 1
57
for nonspontaneous redox reaction(in standard state, the reverse direction)
- deltaGdegree > 0 | - Edegree cell
58
Faradays Constant(F)
the charge in coulombs of 1 mol of electrons
59
deltaGdegree = -nFEdegree cell
the relationship between delta Gdegree and Edegree cell
60
Edegree cell = (0.0592V/n)log(K)
the relationship between Edegree cell and K
61
deltaGdegree = -RT(lnK)
deltaGdegree = -RT(lnK)
62
Nernst Equation
Ecell = Edegree cell – (0.0592V/n)logQ
63
Batteries are
voltaic(galvanic) cells
64
Fuel Cell
the reactants(the fuel provided by an outside source) constantly flow through the battery generating electrical current as they undergo a redox reaction - key difference from a battery: not self contained but can continue to be used unlike a battery - Hydrogen-Oxygen fuel cell only creates H2O which is good!
65
Corrosion
usually gradual, nearly always undesired, oxidation of metals that are exposed to oxidizing agents in the environment
66
the oxidation(or corrosion) of many metals is spontaneous when paired
with the reduction of oxygen
67
anodic regions
where oxidation reactions tend to occur at defects on the surface of a metal(scratch in paint)
68
important conditions about rust:
- moisture must be present(water) - additional electrolytes promote rusting(salt makes car rust faster) - presence of acids promotes rusting(lower pH increases the rate of reduction involving oxygen)
69
preventing rusting:
- no moisture contact - sacrificial electrode – coat the iron in an electrode that oxidizes in irons place - coat iron in a metal that oxidizes more easily(zinc coated nails oxidize and form a strong coating protecting the iron from any future rusting)
70
Electrolysis occurs in
electrolytic cells
71
electrolysis
electrical current drives an otherwise nonspontaneous redox reaction - putting energy in to do work
72
In all electrochemical cells
- oxidation occurs at the anode | - reduction occurs at the cathode
73
In voltaic cells
- the anode is the source of electrons and has the negative charge(anode -) - the cathode draws electrons and has a positive charge(cathode +)
74
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 power course(cathode -)
75
in the electrolysis of a pure molten salt
the anion is oxidized and the cation is reduced
76
the cation that is most
easily reduced(the one with more positive electrode potential) is reduced first
77
the anion that is most
easily oxidized(the one with the more negative electrode potential) is oxidized first
78
the half reaction that occurs the easiest will
occur when two or more reactions can happen(an ion or water etc)
79
the cations of active metals – those that are not easily reduced(Li^+, K^+, Na^+,Mg^2+,Ca^2+, and Al^3+) cannot be
reduced from aqueous solutions by electrolysis because water is reduced at a lower voltage
80
overvoltage
an additional voltage that must be applied in order to get some nonspontaneous reactions to occur
81
you can use the magnitude of electrons(3 mol e- etc) as a conversion with mols.
(1mol e-/96485C)(1mol Au/3mol e-)

Chem 2 Heider (9 decks)

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