14.1 Flashcards

1
Q

what is the standard electrode potential

A

measuring how easily a substance loses electrons

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2
Q

what is the absolute potential difference

A

potential difference between the metal and the solution

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3
Q

why is it not possible to measure directly the absolute potential difference and how can it be solved

A

because in order to create a circuit there must be another metal dipped in the solution which will interfere with the results
instead a reference electrode is used to measure the potential difference between the metal electrode and the reference electrode

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4
Q

what is the reference electrode

A

standard hydrogen electrode

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5
Q

describe the standard hydrogen electrode

A

hydrogen gas at 100kPa bubbled over platinum foil dipped in 1mol/cm^3 HCl solution at 298k

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6
Q

why is the platinum foil for the standard hydrogen electrode covered in porous platinum

A

it has a large surface area and allows equilibrium between hydrogen ions in solution and hydrogen gas to be reached quickly

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7
Q

what are the standard conditions for electrodes

A

100kPa
298k
1mol/dm^3 solution

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8
Q

what is a salt bridge needed for in a complete cell

A

completes the electrical circuit by allowing the movement of ions

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9
Q

what creates a salt bridge

A

paper soaked in a concentrated solution of potassium nitrate (or any ionic salt which will not interfere with the components of the half cell)

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10
Q

how do electrons move in positive electrodes

A

electrons flow from the hydrogen electrode to the metal electrode

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11
Q

how do electrons move in a negative electrode

A

electrons flow from the metal electrode to the hydrogen electrode

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12
Q

where does the position of equilibrium lie when the E^o value is more negative and is the species a more powerful reducing or oxidizing agent

A

to the left
reducing agent

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13
Q

where does the position of equilibrium lie when the E^o value is more positive and is the species a more powerful reducing or oxidizing agent

A

to the right
oxidizing agent

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14
Q

what is the electromotive force (emf)

A

the standard electrode potential of a half cell measured under standard conditions of 298K, 100kPa, 1mol/dm^3 connected to a standard hydrogen electrode

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15
Q

define standard electrode potential

A

the standard electrode potential of a half cell is the emf of a cell containing the half-cell connected to the standard hydrogen electrode under standard states

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16
Q

what is the cell diagram for the zinc and copper cells

A

Zn(s)/Zn^2+(aq) :: Cu^2+(aq)/ Cu(s)

17
Q

what does :: represent in a cell diagram

A

the salt bridge

18
Q

how do electrons flow between half cells

A

from the more negative E^o value half cell to the less negative E^o value half cell

19
Q

is this reaction thermodynamically feasible: Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)
Zn^2+(aq) + 2e- <=> Zn(s) E^o = -0.76 V
Cu^2+(aq) + 2e- <=> Cu(s) E^o = +0.34 V

A

yes because Zn has a more negative E^o value therefor equilibrium lies to the left and is releasing electrons which are accepted by Cu as it has a more positive E^o value therefor equilibrium lies to the right

20
Q

is this reaction thermodynamically feasible: Zn(s) + 2H+(aq) -> Zn^2+(aq)
Zn^2+(aq) + 2e- <=> Zn(s) E^o = -0.76 V
2H+(aq) + 2e- <=> H2(g) E^o = 0.00 V

A

yes because Zn has a more negative E^o value therefor equilibrium lies to the left and is releasing electrons which are accepted by H as it has a more positive E^o value therefor equilibrium lies to the right

21
Q

is this reaction thermodynamically feasible: Cu(s) + 2H+(aq) -> Cu^2+(aq) + H2(g)
Cu^2+(aq) + 2e- <=> Cu(s) E^o = +0.34 V
2H+(aq) + 2e- <=> H2(g) E^o = 0.00 V

A

no because Cu has a more positive E^o value than H so it is not thermodynamically feasible as the two equilibria move in opposite directions to those required for the reaction

22
Q

is this reaction thermodynamically feasible: Cu^2+(aq) + H2(g) -> Cu(s) + 2H+(aq)
Cu^2+(aq) + 2e- <=> Cu(s) E^o = +0.34 V
2H+(aq) + 2e- <=> H2(g) E^o = 0.00 V

A

yes because H has a more negative E^o value therefor equilibrium lies to the left and is releasing electrons which are accepted by Cu as it has a more positive E^o value therefor equilibrium lies to the right

23
Q

is this reaction thermodynamically feasible:
MnO2(s) + 4HCl(aq) -> Mn^2+(aq) + 2Cl-(aq) + 2H2O(l) + Cl2(g)
MnO2(s) + 4H+(aq) + 2e- <=> Mn^2+(aq) + 2H2O(l) E^o = +1.23
Cl2(g) + 2e- <=> 2Cl-(aq) E^o = +1.36

A

no because Cl has a more positive E^o cell value than MnO2 so it is not thermodynamically feasible as the two equilibria move in opposite directions to those required for the reaction

24
Q

How can you make this reaction thermodynamically feasible
MnO2(s) + 4HCl(aq) -> Mn^2+(aq) + 2Cl-(aq) + 2H2O(l) + Cl2(g)
MnO2(s) + 4H+(aq) + 2e- <=> Mn^2+(aq) + 2H2O(l) E^o = +1.23
Cl2(g) + 2e- <=> 2Cl-(aq) E^o = +1.36

A

increasing the concentration of HCl will shift the position of equilibrium to the left and the E^o value becomes less positive and the reaction becomes thermodynamically feasible (not under standard conditions)

25
why may a reaction not occur even if the standard electrode potentials indicate the reaction is thermodynamically feasible
the activation energy may be too large the reaction may not be taking place under standard conditions
26
explain the disproportionation of copper using electrode potentials
Cu^2+(aq) + e- <=> Cu+(aq) E^o = +0.15V Cu+(aq) + e- <=> Cu(s) E^o = +0.52V the first equilibrium is more negative therefor it goes to the left (oxidation) and the second equilibrium is more positive therefor it goes to the right (reduction)
27
what is the equation relating #G and E^ocell
#G = -nFE^ocell n - number of moles involved in reaction F - faraday constant
28
what is the equation relating lnK and E^ocell
lnK = nFE^ocell/RT