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Chemistry A Level > 1.11 Electrode Potentials > Flashcards

1.11 Electrode Potentials Flashcards

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1
Q

Give the equilibrium formed when you dip zinc in a solution of ions

A

Zn2+ (aq) + 2e- <=> Zn(s)

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2
Q

Which side of the equilibrium
Zn2+ + 2e- <=> Zn
Is negative / positive

A

Left has a negative potential due to the build up of electrons, and right is positive as electrons are used up

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3
Q

The larger a potential difference, the more reactive or unreactive a metal is?

A

The more reactive.

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4
Q

What are the three types of electrodes?

A
  • Metal electrodes
  • Gas electrodes
  • Redox electrodes
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5
Q

What is a metal electrode?

A

It consists of a metal surrounded by a solution of its ions eg
Zn(s) | Zn2+(aq)

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6
Q

What is a gas electrode?

A

Consists of a gas and a solution of its ions. An inert metal (usually platinum) is used as the electrode to allow the flow of electrons
e.g. Pt(s) | H2(g) | H+(aq)

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7
Q

What is a redox electrode?

A 
Consists of two different ions of the same element in solution with an inert metal electrode to allow the flow of electrons 
Eg Pt(s) | Fe2+(aq),Fe3+(aq)
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8
Q

What is an electrochemical cell?

A

It is formed when two electrodes in electrolytes are connected, resulting in the generation of an electric current which is caused by the different half equations

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9
Q

What is the difference in electrode potential between the two potentials in an electrochemical cell?

A

The electromotive force (emf) and is a measure of the force which moves the electrons around the circuit.

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10
Q

What is a salt bridge?

A

A paper soaked in potassium chloride or nitrate which completes the circuit

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11
Q

How does a salt bridge complete the circuit

A

It acts as a conductor between the two half-reactions which have to be kept separate, i.e. by being an ionic connection between two ionic solutions in which the ions are free to move and conduct the charge.

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12
Q

What are the standard conditions of a half cell?

A
  • Cell concentration of 1.0 Moldm-3 of the ions involved in the half equation
  • Cell temperature of 298K
  • Cell pressure of 100kPa
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13
Q

What is the standard hydrogen electrode?

A

Potentials of all electrodes are measured by comparing their potential to that of the standard hydrogen electrode. Called the primary standard as it is the standard to which all other potentials are compared.

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14
Q

Give the cell notation of the standard hydrogen electrode.

A

Pt(s) | H2(g) | H+(aq)

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15
Q

What voltage is the standard hydrogen electrode?

A

It is given a value of 0V by definition

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16
Q

Is the cathode positive or negative, and is it where reduction or oxidation occurs?

A

Reduction occurs here, and it is positive

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17
Q

Is the anode positive or negative, and is it where reduction or oxidation occurs?

A

Oxidation occurs here, and it is negative

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18
Q

What does a single | represent in a cell?

A

It indicates a phase change

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19
Q

What does a || represent in a cell?

A

A salt bridge.

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20
Q

What are written next to the salt bridge in cell notation?

A

The most oxidised aqueous ion.

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21
Q

How are the half cells in a cell notation arranged?

A

The half cell with the more positive potential is set up as the right hand electrode, except when measuring potentials against primary standard electrodes when the standard electrode is always the left hand electrode.

22
Q

How do you calculate the emf of a cell?

A

E°cell = E°R - E°L

E°cell = E°Reduction - E°Oxidation

23
Q

Why is the standard hydrogen electrode not ideal?

A
  • It is difficult to set up
  • H2 is very explosive
  • Standard conditions are very hard to maintain
  • It is time consuming to use.
24
Q

Give an example of a secondary standard that is used instead of the SHE.

A

Calomel electrode :
HgCl2 + 2e- <=> 2Hg + 2Cl-
AgCl2 + e- <=> Ag + 2C,-

25
Q

What is the electrochemical series?

A

a list of electrode potentials in order of decreasing (or increasing) potential.

26
Q

Which electrode gains electrons (and is reduced)

A

The more positive electrode gains electrons

27
Q

What causes potential difference between half-cells ?

A

If the position of equilibrium in one of the half-cells lies further to the left than that of the other half-cell, the electrodes have a potential difference.

28
Q

What cells will occur spontaneously (and therefore work)

A

Only a cell with a positive E°cell value (voltage)

29
Q

What are the three main types of commercial cells?

A
  • Non-rechargeable
  • Rechargeable
  • Fuel cells
30
Q

What is an advantage and a disadvantage of commercial cells?

A

+Portable source of electricity

-Creates waste

31
Q

Describe non-chargeable commercial cells.

A

Chemicals are used up over time and emf drops until one chemical is completely used up and the emf = 0V.
Cannot be recharged and must be disposed of after their use.

32
Q

Give the equation for the reaction occurring in a zinc-carbon cell.

A

2MnO2 + 2H+ + Zn + 2NH3 -> Mn2O3 + H2O + Zn(NH3)2^2+

33
Q

Give the cell notation of a zinc-carbon commercial cell.

A

Zn(s) | Zn(NH3)2^2+ (aq) || MnO2(s), Mn2O3(s) | C(s)

34
Q

Give the equation for the reaction occurring in an alkaline non-rechargeable cell.

A

2MnO2 + 2H2O + Zn -> 2MnO(OH) + 2Zn(OH)2

35
Q

Give the cell notation for the alkaline non-rechargeable cell.

A

Zn(s) | Zn2+(aq) || MnO2(s),MnO(OH)(s) |C(s)

36
Q

Describe rechargeable commercial cells.

A

Reactions are reversible and are reversed by applying an external current which regenerates the chemicals.

37
Q

Give the equation for a lithium rechargeable cell, with the forward reaction being in use and backwards being recharging.

A

Li + CoO2 <=> LiCoO2

38
Q

Give the cell notation for the lithium ion rechargeable cell.

A

Li(s) | Li+(aq) ||CoO2(s)| Li+(aq), CoO2^- (aq) | C(s)

39
Q

Give the equation for a lead-acid rechargeable cell, with the forward reaction being in use and backwards being recharging.

A

PbO2(s) + 2H2SO4(aq)+Pb(s) <=> 2PbSO4(s)+2H2O(l)

40
Q

Give the cell notation for the lead-acid ion rechargeable cell.

A

Pb(s)|PbSO4(s)|H+(aq),HSO4-(aq)||PbO2(s),PbSO4(s)|Pb(s)

41
Q

Give the equation for a nickel-cadmium rechargeable cell, with the forward reaction being in use and backwards being recharging.

A

NiO(OH)+2H2O+Cd <=> Ni(OH)2 + Cd(OH)2

42
Q

Give the cell notation for the nickel-cadmium ion rechargeable cell.

A

Cd(s)|Cd(OH)2(s)||NiO(OH)(s), Ni(OH)2(s) | Ni(s)

43
Q

What are fuel cells?

A
  • Cells that require a continuous supply of chemicals in order to ensure that they do not run out of chemicals or need recharging.
  • Can be run in alkaline or acidic conditions
44
Q

Give the overall equation for a hydrogen fuel cell.

A

2H2(g) + O2(g) -> 2H2O(l)

45
Q

Give the cell notation for an acidic hydrogen fuel cell.

A

Pt(s)|H2(g)|H+(aq)||O2(g)|H2O(l)|Pt(s)

46
Q

Give the cell notation for an alkaline hydrogen fuel cell.

A

Pt(s)|H2(g)|H2O(l)||O2(g)|OH-(aq)|Pt(s)

47
Q

Give the overall equation for an ethanol fuel cell.

A

C2H5OH + 3O2 -> 2CO2 + 3H2O

48
Q

Give the cell notation of an ethanol fuel cell.

A

Pt(s)|C2H5OH(l)|CO2(g)||O2(g)|H2O(l)|Pt(s)

49
Q

Give the overall equation for a methanol fuel cell.

A

CH3OH + 1.5O2 -> CO2 + 2H2O

50
Q

Give the cell notation for a methanol fuel cell.

A

Pt(s)|CH3OH(l)|CO2(g)||O2(g)|H2O(l)|Pt(s)

51
Q

Give some benefits and weaknesses of rechargeable cells.

A

+Less waste, cheaper in the long run, lower environmental impact
-Some waste issues (at end of useful life)

52
Q

Give some benefits and weaknesses of hydrogen fuel cells.

A

+Only waste product is water, do not need recharging, very efficient
-Need constant supply of fuels, hydrogen is flammable and explosive, hydrogen usually made using fossil fuels, high cost of fuel cells

Chemistry A Level (35 decks)

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