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a level chemistry > 3.1 redox and standard electrode potential > Flashcards

3.1 redox and standard electrode potential Flashcards

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1
Q

what does the electrode potential mean?
(dont need definition just understand it)

A

how easily a metal looses its electrons into a solution of its own ions

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2
Q

a combined system containing two half cells is called..?

A

an electrochemical cell

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3
Q

if a half cell is connected to another half cell with a different metal, there will be a potential difference or electromotive force (emf) between the two electrode, causing a current to flow between them

A
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4
Q

how can the e.m.f be measured?

A

using a voltmeter

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5
Q

what is an advantage of using a voltmeter to measure the e.m.f?

A

voltmeters have a high resistance so that they do not divert much current from the main circuit

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6
Q

in what direction do electrons flow in electrochemical cells?

A

from the left to the right

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7
Q

in an electrochemical cell, which side is the negative side?

A

LHS

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8
Q

what is the salt bridge used for in an electrochemical cell?

A

to complete the circuit (without allowing solutions to mix)

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9
Q

the salt bridge is often a piece of filter paper saturated with a solution of an inert (unreactive) electrolyte such as KNO3 (aq)

A
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10
Q
  • in electrochemical cells, the better reducing agents are on the (RHS/LHS)?
  • the better oxidising agents are on the (RHS/LHS)??
A
  • LHS
  • RHS
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11
Q

dont use the names anode and cathode

A
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12
Q

for half equations showing reduction, what side of the equations are the electrons??

A

left (LHS)

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13
Q

if the half-reaction doesnt contain a metal in its elemental state, what must be used?

A

an inert platinum electrode

  • this is required in order to connect the redox couple to the external circuit.

THE Pt ELECTRODE DOESNT GET INVOLVED IN THE REACTION

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14
Q

if a gas is involved in electrochemical cells, what must happen?

A
  • it must be bubbled through the solution in such a way that it is in contact with the electrode
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15
Q

what does the electrode potential depend on?

A
  • the conditions used:
    • temperature
    • pressure
    • concentration of reactants
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16
Q

what are the standard conditions for electrochemical cells?

A
  • 298K
  • a pressure of 1atm
  • all species in solution having a concentration of 1 moldm^-3
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17
Q

electrode potentials measured under standard conditions are known as what?

A
  • standard electrode potentials
  • E°
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18
Q

why do you use an electrode made of platinum when you want to measure the electrode potential of something thats not a metal (e.g liquid)?

A

because its inert + you dont want it to react

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19
Q

the emf of electrochemical cells is easy to measure, but the individual electrode potentials themselves cant actually be measured

A

it is only possible to measure the potential difference between two electrodes

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20
Q

what are all electrode potentials measured relative to?

A

the Standard Hydrogen Electrode (SHE)

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21
Q

how to calculate the emf?

A

EMF = E(RHS) - E(LHS)

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22
Q

the more negative a system is, the better ____ agent it is?

A

reducing

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23
Q

what do E° numbers vary with?

A
  • temperature
  • pressure
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24
Q
  • all half equations for electrochemical cells are written as reduction processes
A
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25
Q

physically drawing an electrochemical cell each time is time consuming, so this process can be summarised as a cell diagram: what are some of the rules to follow:

A
  • two vertical double lines represents the salt bridge in the middle of the electrochemical cell (you may see this as a single vertical dotted line)
  • on the LHS of the salt bridge, the order of placing species is as follows:
    • electrode —> reduced species —> oxidised species
  • on the RHS of the salt bridge, the order of placing species is as follows:
    • oxidised species —> reduced species —> electrode
  • include state symbols when writing these out
  • a change of state between one species and the next is represented by a solid line
  • a non-change of state is indicated by a comma
  • be aware that sometimes your electrode will be the reduced species
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26
Q

if the EMF is (positive/negative), we can say that the reaction is feasible/spontaneous and hence will occur??

A

positive

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27
Q

the more positive the EMF, the more ____ the reaction?

A

feasible

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28
Q

can cell potentials be used to predict how fast a reaction can happen?

A

no
- they CAN be used effectively to predict whether or not a given reaction will take place
- but they CANT give any indication as to how fast a reaction will proceed

29
Q

if something has a positive Ecell but no apparent reaction occurs it’s because the reactants are kinetically stable, has a high activation energy so is very slow at room temperature

A
30
Q

if reactions are expected to occur but dont seem to, what could be the reasons?

A
  • too dilute (conditions not standard)
  • reactions too slow (reactants are kinetically stable)
31
Q

if a reaction is not expected to take place but does, why might this happen?

A
  • because the conditions are non-standard (ie solutions are concentrated)
32
Q

what can be used to predict how electrode potentials are affected when non-standard conditions are used?

A

le chatelier’s principle

33
Q
  • if the oxidizing agent has a concentration greater than 1moldm^-3, it is more likely to favour reduction and the electrode potential will be more positive than the standard electrode potential
  • if it has a concentration of less than 1mol^-3, it is more likely to favour oxidation and the electrode potential will be more negative than the standard electrode potential
  • for reducing agents, the reverse is true
A
34
Q

e.g Fe2+ (aq) + 2e- ⇌ Fe (s)
standard electrode potential = -0.44V
if [Fe2+]= 0.1moldm^-3 the electrode potential = -0.50V

A
  • the concentration is lower than standard so reduction is less likely to take place, and hence the electrode potential is more negative than expected
  • if the temperature is higher than 298K, then the system will move in the endothermic direction and the electrode potential will change accordingly
  • if the pressure is greater than 1atm, then the system will move to decrease the pressure and the electric potential will change accordingly
35
Q
  • a change which favours the reduction direction will make the electrode potential more ____
  • a change which favours the oxidation direction will make the electrode potential more ____
A
  • positive
  • negative
36
Q

e.g Fe2+ + e- ⇌ Fe+
E° = 0.50v
- what happens to E° when decreasing the concentration of Fe2+?

A
  • equilibrium shifts to LHS
  • oxidation is favoured (bc e- lost)
  • ∴ better/more effective reducing agent
  • E° value becomes more negative
37
Q

e.g Cu2+ ⇌ Cu3+ + e-
E° = +0.34V
∆H = +
what happens to the E° if the temperature increases?

A
  • ∆H = + ∴ endo
  • equilibrium shifts to RHS
  • oxidation favoured
  • ∴ more effective reducing agent
  • E° value becomes more negative
38
Q

an increase in temperature favours (endothermic/exothermic) reactions?

A

endothermic

39
Q

what is a fuel cell?

A
  • a cell in which a chemical reaction between a fuel and oxygen is used to create a voltage
  • the fuel and oxygen flow into the cell continuously and the products flow out of the cell
  • therefore the cell does not need to be recharged
40
Q

what is the most widely used fuel cell?

A

hydrogen-oxygen fuel cell

41
Q

what does a fuel cell consist of?

A
  • two half-cells connected by a semi-permeable membrane
  • an aqueous solution of sodium hydroxide is used as the electrolyte
42
Q

hydrogen fuel cell:

A
  • oxygen is pumped into one of the half-cell:
    O2 (g) + H2O (l) + 4e- —> 4OH- (aq) E°=+0.40V
  • hydrogen is pumped into the other half-cell:
    H2O (l) + 2e- —> H2 (g) + 2OH- (aq) E°=-0.83V
  • the oxygen half cell is more positive and therefore undergoes reduction
  • the hydrogen half cell is more negative and undergoes oxidation:
    O2 + H2O + 4e- —> 4OH- reduction
    H2 + 2OH —> H2O + 2e- oxidation

O2 (g) + 2H2 (g) —> 2H2O (l) overall cell reaction, emf=1.23V
- hydroxide ions are generated in the oxygen half cell and travel through the membrane into the hydrogen half cell, where they are used up
- water is the product of the reaction and is allowed to off

43
Q

what are some advantages of using hydrogen fuel cells?

A
  • the hydrogen-oxygen fuel cell produces water as the only product (∴ doesn’t produce any greenhouse or polluting gases associated with combustion energies. the process of generating hydrogen for use in fuel cells produces a small quantity of CO2 but much less than would be generated by a combustion engine
  • fuel cells are more efficient than combustion energies
  • process is continuous as long as the fuel is supplied
44
Q

what are some limitations of using hydrogen fuel cells?

A
  • hydrogen is a flammable gas with a low boiling point - it is ∴ both difficult and dangerous to store and transport. It can be stored as a liquid under pressure or as a solid adsorbed to the surface of a solid, but both of these techniques are expensive
  • as a result obtaining hydrogen as a fuel is difficult - and this means that people will not buy hydrogen powered vehicles
  • fuel cells use toxic chemicals in their manufacture
  • fuel cells have a limited lifetime
  • efficiency is affected by temperature

(-hydrogen is expensive and hard to store
- high pressure tanks are required to store oxygen and fuels like hydrogen
- materials used to make them are expensive)

45
Q

OIL
RIG

A
46
Q

what is a redox reaction?

A
  • a reaction in which oxidation and reduction occur on different species simultaneously
47
Q

what is standard electrode potential?

A
  • the potential across the electrodes when a redox system is connected to a hydrogen half-cell under standard conditions
48
Q

what is the standard electrode potential for hydrogen assumed to be ?
how is this an advantage to it being the standard electrode potential?

A
  • 0 volts
  • so it is used to allow for easy comparison between thr electrode potential of different elements
49
Q

what is the experimental setup to calculate the standard electrode potential for zinc?

A

LEFT:
- hydrogen standard cell always placed on left
- Pt electrode
- H2 enters
- solution of HNO3 - dissociation into H+

  • e.g NaCl salt bridge

RIGHT:
- zinc electrode
- Zn(NO3)2 solution - dissociation into Zn2+

50
Q

why must metal electrodes be cleaned with sandpaper before creating an electrochemical cell?

A
  • to remove any metal oxide that has formed on the surface and improve electrical conductivity
51
Q

describe the movement of electrons in an electrochemical cell

A
  • electrons flow through the wire from the positive electrode to the negative electrode
52
Q

why must an inert salt bridge be used in the salt bridge?

A
  • the salt must be inert so that it doesn’t react with the solutions and alter the ion concentrations
  • if a reactive salt was used, the cell potential would change
53
Q

what moves across the salt bridge?

A

ions

54
Q

for what range of cell potential values is a process feasible?

A
  • cell potential must be greater than 0
55
Q

why might theoretical cell potential values be different to values obtained experimentally?

A
  • conditions may be non-standard
56
Q
  • a cell is made up of the following half cells:
    Ag+ (aq) + e- ⇌ Ag (s) E°=+0.80V
    Cu2+ (aq) + 2e- ⇌ Cu (s) E°=+0.34V
  • write the overall cell equation and calculate the standard cell potential
A

2Ag + (aq) + Cu(s) —> Ag(s) + Cu2+ (aq)
E°cell= +0.80 - (+0.34) = 0.46V

57
Q

in an electrochemical cell, is the more negative half cell oxidised or reduced?

A

oxidised

58
Q

what is the only product of a hydrogen-oxygen fuel cell?

A

water

59
Q

how does a hydrogen-oxygen fuel cell work?

A
  • hydrogen and oxygen are pumped through porous electrodes. the electrolyte is often an acid such as phosphoric acid
  • hydrogen and oxygen react, producing energy and water
60
Q

what are the 2 half equations taking place in a hydrogen-oxygen fuel cell?

A

2H2 + 4OH- ⇌ 4H2O + 4e-
O2 + 2H2O + 4e- ⇌ 4OH-

61
Q

write an equation for the overall reaction that takes place in a hydrogen-oxygen fuel cell

A

2H2 + O2 —> 2H2O

62
Q

the gas oxygen, O2, is converted into ozone, O3, in the upper atmosphere. the equation for this process is:
3O2 —> 2O3
use oxidation states to explain why this is not a redox reaction [2]

A
  • both O2 and O3 have oxidation states of zero
  • no change in oxidation state
63
Q
  • an important technological development in recent years has been the hydrogen fuel cell
  • this uses electrochemical methods to get energy from hydrogen
  • write the half-equations for the processes occurring at the electrodes and an equation for the overall reaction [3]
A
  • (at anode) H2 —> 2H+ + 2e-
  • (at cathode) O2 + 4H+ + 4e- —> 2H2O
  • (overall) 2H2 + O2 —> 2H2O
64
Q
  • the following equations show the standard electrode potential for the Cl2/Cl-and I2/I- systems
  • Cl2 + 2e- ⇌ 2Cl- E°= +1.36V
  • I2 + 2e- ⇌ 2I- E°= +0.54V
  • use these values to explain why only hydrogen iodide (represented as I- in the equation) is able to further react with concentrated sulphuric acid in this way [2]
A
  • the values show that chlorine is the best oxidising agent, as it has the most positive E° value and therefore iodide is the better reducing agent
  • and is strong enough to reduce the sulfuic acid
65
Q

a more positive E° value means it is a stronger ____ agent?

A

oxidising

66
Q

define the term standard electrode potential [3]

A
  • emf/potential difference/voltage (of electrochemical cell)
  • between a standard hydrogen electrode and a half cell
  • 1 atm, 1moldm^-3, 298k
67
Q

consider the half reaction:
Fe2+ + 2e- ⇌ Fe
define the term standard electrode potential with reference to this electrode [3]

A
  • emf of cell / potential difference of cell containing Fe2+ AND Fe
  • AND standard hydrogen electrode
  • 1 moldm^-3, 1 atm, 298k
68
Q

explain how the salt bridge provides an electrical connection between the two solutions [1]

A
  • completes the circuit
  • without allowing solutions to mix

(it allows ions to move through it)

a level chemistry (17 decks)

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