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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
Q

in a voltaic cell the potential difference between the two electrodes is

A

the cell potential(Ecell) or cell emf

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26
Q

cell potential depends on

A

the relative tendencies of the reactants to undergo oxidation and reduction
- stronger the tendencies the larger the difference and therefore a higher potential

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27
Q

cell potential also depends on

A

the concentrations of the reactants and products and temperature

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28
Q

standard cell potential(Edegree symbol cell) or standard emf

A

the cell potential at standard conditions(1 atm for gaseous reactants and 1M for solutions)

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29
Q

cell potential is a measure of

A

the overall tendency of the redox reaction to occur spontaneously

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30
Q

a negative cell potential indicates

A

that the forward reaction is not spontaneous

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31
Q

anode

A

in all electrochemical cells the electrode where oxidation occurs

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32
Q

cathode

A

in all electrochemical cells the electrode where reduction occurs

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33
Q

in a voltaic cell the anode is

A

the more negatively charged electrode labeled with the negative(-) sign

34
Q

in a voltaic cell the cathode is

A

the more positively charged electrode and labeled with the positive(+) sign

35
Q

electrons flow from the anode to cathode through wires

A

from negative to positive

36
Q

as electrons flow out of the anode positive charge builds up in the solution

A

form in the oxidation half cell

37
Q

as electrons flow into the cathode positive ions are reduced at the reduction half cell

A

negative charge builds up in the solution

38
Q

salt bridge

A

an inverted U shaped tube that contains a strong electrolyte such as KNO3 and connects the two half cells

39
Q

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

A
  • the salt bridge completes the circuit allowing electrical current to flow
40
Q

standard cell potential(Edegree sign cell for an electrochemical cell depends on

A

the specific half reactions occurs and the potential energy difference(per unit charge) of the two electrodes

41
Q

standard electrode potential

A

each electrode in each half cell has its own individual potential

42
Q

the overall cell potential(Edegree Cell) is

A

the difference between the two standard electrode potentials

43
Q

while cells are connected electrons flow from the electrode with

A

more negative charge(greater potential energy) to the electrode with more positive charge(less potential energy)

44
Q

standard hydrogen electrode(SHE)

A

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
Q

Edegree cell =

A

Edegree final – Edegree initial

- Edegree cathode – Edegree anode

46
Q

the electrode potential of the standard hydrogen electrode(SHE) =

A

0

47
Q

the electrode in any half-cell with a greater tendency to undergo reduction is

A

positively charged relative to the SHE and therefore has a positive Edegree cell

48
Q

the electrode in any half-cell with a lesser tendency to undergo reduction(or greater tendency to undergo oxidation) is

A

negatively charged relative to the SHE and therefore has a negative Edegree cell

49
Q

the cell potential of any electrochemical cell(Edegree cell) is the difference between

A

the electrode potentials of the cathode and anode(Edegree cell = Edegree cathode – Edegree anode)

50
Q

Edegree cell is positive for _______ reactions and negative for ______ reactions

A

spontaneous

nonspontaneous

51
Q

NIO

A

more negative is oxidation

52
Q

PIR

A

more positive is reduction

53
Q

the half reaction with the more positive electrode potential

A

attracts electrons more strongly and will undergo reduction

54
Q

the half reaction with the more negative electrode potential

A

repels electrons more strongly and will undergo oxidation

55
Q

any reaction is table 18.1 is spontaneous when paired with

A

the reverse of any of the reactions listed below it on the table

56
Q

For a spontaneous redox reaction(in standard states the forward direction)

A
  • deltaGdegree 0

- K > 1

57
Q

for nonspontaneous redox reaction(in standard state, the reverse direction)

A
  • deltaGdegree > 0

- Edegree cell

58
Q

Faradays Constant(F)

A

the charge in coulombs of 1 mol of electrons

59
Q

deltaGdegree = -nFEdegree cell

A

the relationship between delta Gdegree and Edegree cell

60
Q

Edegree cell = (0.0592V/n)log(K)

A

the relationship between Edegree cell and K

61
Q

deltaGdegree = -RT(lnK)

A

deltaGdegree = -RT(lnK)

62
Q

Nernst Equation

A

Ecell = Edegree cell – (0.0592V/n)logQ

63
Q

Batteries are

A

voltaic(galvanic) cells

64
Q

Fuel Cell

A

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
Q

Corrosion

A

usually gradual, nearly always undesired, oxidation of metals that are exposed to oxidizing agents in the environment

66
Q

the oxidation(or corrosion) of many metals is spontaneous when paired

A

with the reduction of oxygen

67
Q

anodic regions

A

where oxidation reactions tend to occur at defects on the surface of a metal(scratch in paint)

68
Q

important conditions about rust:

A
  • 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
Q

preventing rusting:

A
  • 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
Q

Electrolysis occurs in

A

electrolytic cells

71
Q

electrolysis

A

electrical current drives an otherwise nonspontaneous redox reaction
- putting energy in to do work

72
Q

In all electrochemical cells

A
  • oxidation occurs at the anode

- reduction occurs at the cathode

73
Q

In voltaic cells

A
  • the anode is the source of electrons and has the negative charge(anode -)
    - the cathode draws electrons and has a positive charge(cathode +)
74
Q

In electrolytic cells

A
  • 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
Q

in the electrolysis of a pure molten salt

A

the anion is oxidized and the cation is reduced

76
Q

the cation that is most

A

easily reduced(the one with more positive electrode potential) is reduced first

77
Q

the anion that is most

A

easily oxidized(the one with the more negative electrode potential) is oxidized first

78
Q

the half reaction that occurs the easiest will

A

occur when two or more reactions can happen(an ion or water etc)

79
Q

the cations of active metals – those that are not easily reduced(Li^+, K^+, Na^+,Mg^2+,Ca^2+, and Al^3+) cannot be

A

reduced from aqueous solutions by electrolysis because water is reduced at a lower voltage

80
Q

overvoltage

A

an additional voltage that must be applied in order to get some nonspontaneous reactions to occur

81
Q

you can use the magnitude of electrons(3 mol e- etc) as a conversion with mols.

A

(1mol e-/96485C)(1mol Au/3mol e-)

Chem 2 Heider (9 decks)

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