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electrochemistry > book 2!! > Flashcards

book 2!! Flashcards

1
Q

what is an electrolyte

A

a substance that, when dissolved, gives a solution that conducts electricity

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2
Q

what type of substance are electrolytes

A

usually ionic substances

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3
Q

examples of electrolytes aka ionic substances

A

NaCl
KCl
Na2SO4

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4
Q

in an electrolyte solution,, how is current carried between electrodes

A

carried by the dissolved ions

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5
Q

what metal is used for electrodes

A

platinum

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6
Q

why is platinum used as the metal in a cell for electrodes

A

it’s inert
it’s unreactive

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7
Q

what is applied between the electrodes

A

voltage is applied
resistance is measured (R)

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8
Q

what is the resistance a measure of

A

it’s a measure of how difficult the motion of ions is between electrodes

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9
Q

what are the similarities and differences between the 2 electrodes

A

both platinum
same surface area
one positive (attracts negative ions)
one negative (attracts positive ions)

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10
Q

what does conductivity depend on

A

the conc
the higher the conc,, aka the higher the number of ions in solution,, aka more charge carriers available,, aka higher conductivity

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11
Q

strong electrolytes

A

fully ionise in solution
KCl
Na2SO4

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12
Q

when can ions behave independently in solution

A

when the solution is very dilute
when the ions are far apart

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13
Q

what’s an anion

A

negative ion

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14
Q

what’s a cation

A

positive ion

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15
Q

use of lithium instead of zinc does what to a battery

A

reduces its weight
gives a higher voltage (lithium is more reactive than zinc)

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16
Q

negative about alkaline and lithium batteries

A

they cannot be recharged.
their redox reactions aren’t reversible,, once the reaction has reached equilibrium,, it’s dead

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17
Q

how can batteries be recharged

A

by passing an electric current through them and so reversing the redox reaction

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18
Q

what does recharging a cell require

A

it requires the use of electricity to drive a chemical reaction in the non spontaneous direction

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19
Q

how is charge carried between the electrodes

A

by the dissolved ions from the electrolyte

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20
Q

what is an electrolyte

A

a substance that,, when dissolved,, gives a solution that conducts electricity

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21
Q

why is platinum used as an electrode

A

it’s inert

very unreactive

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22
Q

what is applied on the electrodes

A

voltage!!

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23
Q

what is measured when voltage is applied on the electrodes

A

the resistance is measured
R

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24
Q

what is resistance

A

a measure of how difficult the motion of ions is between electrodes

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25
Q

positive ions move towards what

A

they move towards the negative electrode

the anode

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26
Q

negative ions move towards the

A

the cathode

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27
Q

what happens at the anode

A

oxidation

the negative electrode loses electrons to give to the positive ions

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28
Q

what happens at the cathode

A

reduction

the cathode gains electrons from the negative ions

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29
Q

what does conductivity depend on

A

the number of ions in solution

a higher concentration of ions = more charge carriers available,, higher conductivity

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30
Q

strong electrolytes areeee

A

fully ionised in solution

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31
Q

strong electrolytes + conductivity

A

dilute = higher conductivity

due to interactions between ions affecting the movement of ions through the solution.

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32
Q

describe a daniel cell

A

Zinc in ZnSO4 LHS
Copper in CuSO4 RHS

1 moldm-3 each

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33
Q

what is current

A

the flow of electrons between the half cells

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34
Q

what is a salt bridge used for

A

charge is conducted by ions in solution. current is carried by ions

soaked in sodium sulfate

keeps electrical neutrality

flow of ions between cells

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35
Q

LHS is for the

A

anode

oxidation occurs here

Zn(s) -> Zn2+

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36
Q

RHS is for the

A

cathode

reduction occurs here

Cu2+ > Cu (s)

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37
Q

cell diagram steps

A

X (s) | X (aq) || Y (aq) | Y (s)

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38
Q

what is | in a cell diagram

A

phase boundary

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39
Q

what is || in a cell diagram

A

salt bridge

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40
Q

can we measure the potential of a single half cell

A

nope!!

the volt meter shows the potential difference between the 2 half cells

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41
Q

what can we use to measure the potential of a single half cell

A

use a cell to compare the rest to!!

use a standard hydrogen half cell

42
Q

standard hydrogen half cell equation

A

2 H+ (aq) + 2e- —> H2 (g)

43
Q

potential if a hydrogen half cell + what conditions it needs to be in

A

standard conditions: 1moldm-3, 1bar, 298k

0.00V

44
Q

standard electrode potentials are actually

A

standard reduction potentials

E* values

45
Q

the more positive the E* (reduction potential) the…

A

stronger the oxidising agent (it helps other lose e- by taking them)

the weaker the reducing agent (weaker bc it doesn’t help atoms gain electrons as it’s positive)

46
Q

if the half cells E* value is more positive what does this mean

A
  • oxidising agent
  • accepts electrons
  • reactions are written as reduction : the reaction is more likely to occur in the forward direction (the reduction direction)
47
Q

what if the E* cell value is more negative

A
  • reducing agent
  • loses electrons
  • more likely to occur in the reverse direction : bc reactions are written as reductions but it wants to lose electrons!!
48
Q

the more positive reaction will do what do the more negative reaction

A

more positive,, wants to gain e-

so it oxidises the more negative reaction

49
Q

Br2 added to NaCl and NaI (both 1moldm-3) what happens

A
  • write the reduction equation (electrochem series)
  • Br is added,, so we compare Br with I and Cl
  • Br is more negative than Cl- (it can’t oxidise it - take it’s electrons away)
  • Br is more positive than I- (it can oxidise it,, take its electrons away)

bc I is more negative,, its half reaction will be an oxidation,, we reverse the electrochem reaction) then we combine the 2!! 💗💗

50
Q

equation for finding the E* of the cell!!

A

E* cell = E* cathode - E* anode

E* cell = E* reduction - E* oxidation

51
Q

for the cell reaction to occur spontaneously,, what value must E*cell have

A

A POSITIVE VALUE!!!

52
Q

what if the E* cell isn’t positive

A

the current reaction isn’t spontaneous

it would occur spontaneously in the opposite direction to the cell diagram

53
Q

what does the E*cell value tell us

A

tendency of the overall cell reaction to happen

54
Q

what is an example of the electrochemical cell doing work

A

work = the movement of electrons around the circuit

55
Q

cell doing work!!

A

movement of charge (electrons) around a circuit against resistance.

required energy
work done by the cell.

56
Q

what does the energy required to transfer electrons//charge represent

A

the electrical work the cell can perform

57
Q

what is z

A

moles of electrons

58
Q

what is Q

A

total charge

59
Q

what’s is F

A

faraday constant

charge on 1 mole of electrons

96 485 coulombs

60
Q

total charge produced equation

A

Q = zF

Q= total charge
z= moles
F= faraday constant

61
Q

Electrical work equation

A

-Q x V

electrical work = -charge x potential difference

62
Q

why is electrical work negative

A

bc it takes energy for it to occur

the reaction system loses energy

63
Q

Welectrical can also be =

A
  • zF x E*cell

MAX ELECTRICAL WORK THAT CAN BE TRANSFERRED

64
Q

when does cell voltage = E*cell

A

resistance in circuit is large
zero current flows

65
Q

what must occur for the cell to do work

A

resistance must be lowered so current can flow

66
Q

change in gibbs energy is equal to whattttt

A

-zF x E*cell

equal to the max amount of work a cell can do

G = -zF x E*cell

67
Q

considering gibbs and cell potential,, what values must they have and how are they linked

A

Wele = -zF x E*cell

E*cell = spontaneous if positive

G = -zF x E*cell
gibbs = max work a cell can do

gibbs: spontaneous if negative

68
Q

what conditions is maximum work done in

A

done in reversible conditions

69
Q

what are reversible conditions when it comes to cells

A

high resistance

little current

measured Voltage = E*cell

70
Q

faraday constant units

A

C mol-1

71
Q

E*cell units

A

V

72
Q

z units

A

no units

73
Q

so G = -zF x E*cell units

A

CVmol-1
CV = J

so G = Jmol-1

74
Q

spontaneous Ecell must be

A

positive

75
Q

spontaneous reaction must be

A

negative gibbs

76
Q

gibbs and thermodynamic equilibrium are related in which equation,, and what does this tell us

A

G = -RT ln(K)

tells us we can use the E*cell to find equilibrium constant, K

77
Q

how can we use E*cell to find equilibrium constant K

A

G = -zF x E*cell
G = -RT ln(k)

-zF x Ecell = - RT ln(k)
Ecell = RT/zF ln(k)

ln(K) = zF/RT E*cell

78
Q

what happens when a cell reaction reaches equilibrium

A

it goes flat
it runs out

current is drawn from the cell, the half cell reactions occur, concentration of species changes,, then concentration of species stops changing + no current flows.

state where half cell reaction concentrations stop changing + no more current flows.

79
Q

what does a large equ constant, K mean

A

it means that the reaction basically goes to completion,, all solid X forms Xions

all Y ions form Y solids.

there is complete conversion

80
Q

what if the value of k is 1

A

reaction is at equilibrium

81
Q

what if k is less than 1

A

the reaction is spontaneous in the reverse direction

82
Q

E*cell dependence on concentration

A

E*cells are calculated under standard conditions,, they may change with different concentrations or gas pressures

83
Q

how do we tackle the problem of E*cells being affected by conc and pressure

A

consider how Gibbs changes with rea and pro activity

G = G* + RTln(Q)

84
Q

what is Q

A

reaction quotient

product of (activity x product ^stoich) // product of (activity of reactants ^ stoich) r

85
Q

what does product of mean

A

it means u multiply the different things together.

like sum of but with multiplication instead of addition

86
Q

what is the activity of a solid

A

1

87
Q

Q = K when

A

at equilibrium!!

88
Q

what is the activity of a gas, a liquid and a solid

A

gas : p/p*. pressure of gas / 1bar

liquid: conc/conc* conc/ 1moldm-3

solid = 1

89
Q

finding the cell potential when conditions aren’t standard

A

Ecell = E*cell - (RT/zF) x ln(Q)

90
Q

activity for stuff at standard conditions

A

1!!!

91
Q

G = G* + RTln(Q) can be written as

A

-zF x Ecell = -zF x E*cell + RTln(Q)

92
Q

nernst equation

A

Ecell = E*cell - RT/zF x ln(Q)

93
Q

what does the nernst equation allow us to do

A

find the Ecell (cell potential) under non standard conditions

find cell potential when changing the concentration of the cell

can also find conc of species involved in an electrochemical cell.

94
Q

what is a concentration cell

A

half cells made up of the same chemical reactions but at different concentrations

95
Q

in electrochemical cells,, the chemical reactions are used to make what

A

used to produce energy

96
Q

what is electrolysis

A

driving a chemical reaction by supplying it with electrical energy

97
Q

input of energy allows what reactions to occur

A

non spontaneous reactions occur

98
Q

what cells include reversible reactions that can be recharged

A

secondary cells

99
Q

which battery is used in motor vehicles

A

lead - acid battery

100
Q

electrolysis cell

A

looks like 1 half cell but is connected to an electrical power source (battery)

electrochemistry (11 decks)

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