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Chemistry > Electrode Potential > Flashcards

Electrode Potential Flashcards

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1
Q

How to tell the difference between a reducing agent and an oxidising agent

A

Reducing agents are oxidised, they donate and lose electrons
Oxidising agents are reduced, they accept and gain electrons

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2
Q

How would you define a half cell

A

A strip of metal dipped into a solution of its own ions can be called an electrode or a half cell- when the strip of metal is dipped in solution of its an ions an equilibrium is set up

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3
Q

Using Cu+ as an example, explore how the equilibrium is reached

A
  • the Cu2+ ions dissolve into the solution, giving the solution a positive charge
  • the electrons collect in the copper strip, giving it a negative charge
    -this means a potential difference is established between the two
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4
Q

How does a voltage reading identify where the equilibrium lies

A

If there is a large voltage, the equilibrium is to the right
If there is a small voltage, the equilibrium is to the left

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5
Q

Describe how you would connect two half cells

A

The two metal rods are connected with wires and a high resistance voltmeter
The two beakers of electrolyte are connected with a salt bridge to complete the circuit. This is a glass tube or porous material soaked in potassium nitrate solution

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6
Q

Why is KNO3 a suitable solution for a salt bridge

A

KNO3 is unreactive with the electrodes and the ions are free to move

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7
Q

What is the role of the voltmeter and what happens if it was removed or replaced

A

Prevents electrons flowing and enables the voltage to be measured, if the voltage meter was removed electrons can flow from the left electrode to the right
If the voltmeter is replaced with an ammeter or a bulb, the electrons can flow and a current is produced

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8
Q

Why might the current produced by a cell fall to 0 after some time

A

All the reactants are used up

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9
Q

What will happen to a cell once the reactants are used up

A

Stops working or starts to leak

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10
Q

Why is platinum a suitable electrode

A

Pt is unreactive and conducts electricity

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11
Q

Key points for SHE

A
  1. H2 gas is pumped in at a pressure of 100kPa
  2. The electrolyte contains H+ ions of concentration 1moldm-3
  3. There must be a platinum electrode
  4. The whole system must have a temperature of 298K
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12
Q

What does it mean that the SHE is defined as 0

A

If another electrode is connected up to the standard hydrogen electrode and voltage between them is measured

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13
Q

The standard electrode potential of Cu2+/Cu is 0.37 V. Why might the electrode potential of the following cell not be 0.37V

A

The concentration of the CuSO4 solution is not 1moldm-3

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14
Q

Give an example of how to write conventional cell representation using Zinc and Copper

A

Zn(s)I Zn2+(aq)II Cu2+(aq)I Cu(s)

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15
Q

what des the conventional cell representation show?

A

That at the left electrode, Zinc is oxidised to Zn2+ and at the right electrode Cu2+ is reduced to Cu

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16
Q

Key facts for electrode potential

A
  • the right electrode is the positive electrode
  • the right electrode has a more positive standard electrode potential than the left electrode
  • the left electrode is the negative electrode
  • the left electrode has a less positive standard electrode potential than the right electrode
  • vertical solid lines indicate a phase boundary
  • a double vertical line represents the salt bridge
  • the species with the highest oxidation state should be written closest to the salt bridge
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17
Q

on which side is the SHE always written on

A

the left

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18
Q

using a table, how can you identify which is the strongest oxidising agent and which is the weakest reducing agent

A

the substance with the most positive E° is the weakest reducing agent on the right and the strongest oxidising agent on the left

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19
Q

Why will Ag+ ions react with Li

A

the E° Ag+ is greater than E° for Ag+ and Ag+ is a stronger oxidizing agent

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20
Q

what will be observed when Ag+ reacts with Li

A

Solid Ag forms

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21
Q

What happens when Li+ is added to Ag

A

E° Ag is greater than E° Li+ and Li+ cannot oxidise Ag

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22
Q

Identify a chemical that reacts with Cu

A

Ag+ because E° Ag+ is greater than E°Cu

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23
Q

how to calculate EMF

A

Right-left/ reduction-oxidation/ more positive-least positive

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24
Q

Write an equation for the cell discharge of Silver and Lithium

A

Ag+(aq) + Li(s) –> Ag(s) + Li+ (s)

25
Q

write an equation for the cell recharge of silver and lithium

A

Ag(s) + Li+(s) –> Ag+(aq) + Li(s)

26
Q

give an environmental advantage of using rechargeable cells

A

metals are reused

27
Q

give an environmental disadvantage of using rechargeable cells

A

mains electricity is used to recharge, which may come from combusting fossil fuels, which releases Co2

28
Q

the cell made of the systems Mg/Mg2+ and Fe2+/Fe under standard conditions has an emf of 1.93 V
deduce how the emf changes when the conc of Mg2+ is decreased

Mg2+ + 2e- –> Mg
Fe2+(aq) + 2e- –> Fe

A

Mg2+ electrode equilibrium shifts to the left so E° of Mg2+/Mg gets more negative as more electrons are made and released. Overall EMF increases

29
Q

a diagram shows that the amount of copper in the electrodes is much greater than the amount of copper ions in the copper sulfate solutions- in which direction would the electrons flow in the above cell

A

from right to left
the Cu2+ is more concentrated on the left, so reduction on Cu2+ is more likely to happen on the left
the left electrode is the positive electrode, so the right electrode is the negative electrode

30
Q

single use batteries

A

the electrochemical reaction taking place inside the cell is irreversible- once the reactants have been used up they can’t be used again

31
Q

how is lithium ion battery a rechargeable battery

A

the positive electrode is made of lithium cobalt oxide Li(CoO2) and the negative electrode is made of graphite, when it is discharged, a layer of Li on the negative electrode dissociate to form Li+ and release an electron:
Li—> Li+ + e-
The Li has been oxidised, so this is the negative hand electrode or the left
At the positive electrode a lithium ion and electron combine with CoO2 from the electrolyte to form Li+(CoO2)
CoO2 + Li+ + e- —> Li(CoO2)
Oxidation state is+4 so the Co has been reduced- this is the positive or right hand electrode
Overall equation: CoO2 + Li —> Li(CoO2)

32
Q

What happens when the battery runs out and why does it? Give the equation for recharging

A

Because all the battery has CoO2 or Li has been used up. It can be connected to a stronger electrical current which forces the reaction back in the reverse direction to recharge the battery
Equation for recharging: Li(CoO2)—> CoO2 + Li

33
Q

What is a fuel cell

A

Uses the energy from the reaction of a fuel with oxygen to create a voltage
Modern fuel cells are based off on methanol, natural gas and petrol
Combusting a fuel can can generate a flow of electrons

34
Q

Advantages of using fuel cells for energy instead of fossil fuels

A

Greater efficiency than burning hydrogen in a combustion engine
Less polluting as water is the only product

35
Q

Disadvantages of using fuel cells for energy instead of fossil fuels

A

H2 is difficult to store
Fossil fuels are combusted to produce the hydrogen, which releases CO2

36
Q

Advantages of fuel cells compared to other type of cells

A

Voltage is constant, as fuel and oxygen is supplied constantly so concentrations of reactants remain constant

37
Q

What is an electrolyte

A

An acid or alkaline membrane which allows ions to move from one compartment of the cell to another, almost like a salt bridge

38
Q

Hydrogen oxygen fuel in acidic conditions

A

At the negative electrode: H2 —> 2H+ + 2e-
At the positive electrode: O2 + 4H+ + 4e- —> 2H2O

Overall equation: 2H2 + O2 —> 2H2O

39
Q

Hydrogen oxygen duel cells in alkaline conditions

A

Negative electrode: H2 + 2OH- —> 2H2O + 2e-
Positive electrode: O2 + 2H2O + 4e- —-> 4OH-

Overall equation: 2H2 + O2 —>2H2O

40
Q

Hydrogen oxygen cells in both conditions

A
  • same overall equation, therefore both cells have the same EMF
  • the voltage can remain constant - reagents can be continuously supplied
  • no environmental issues / damage as water is the only product
  • to generate H2 fossil fuels are burnt which produces O2
41
Q

Give one reason why the emf of this cell changes when the electrodes are connected and a current flows

A

The concentration of ions change

42
Q

A student determined the concentration in a solution of iron chloride by titration with acidified potassium dichromate solution. A second student titrated the same solution. Explain why the second student obtained a greater concentration for the iron ions

A

MNO4- will oxidise the chloride ions so a larger volume will be needed

43
Q

Give two reasons why KNO3 is a good reagent for a salt bridge

A

1- doesn’t react
2- ions can move

44
Q

Explain why the ammeter would fall to zero after a time

A

Fe3+ ions would be used up

45
Q

How to identify the weakest oxidising agent

A

It will have the most negative EO value

46
Q

A rechargeable nickel;-cadmium cell is an alternative cell. Deduce the oxidation state of nickel in the cell after recharging is complete
Write an equation for this recharge

A

2Ni(OH)2 + Cd(OH)2 —> 2NiO(OH) + Cd + 2H2O
+3

47
Q

One environmental advantage of rechargeable cells

A

Metal compounds are reused

48
Q

Potassium manganate is an oxidising agent
Suggest why 0.0200 moles-3 of potassium manganate does not need to be kept away from flammable material

A

The solution is very dilute

49
Q

Explain why rechargeable cells are connected to these solar cells

A

Solar cells do not supply electrical energy all of the time- rechargeable cells can store electrical energy for use when the solar cells are not working

50
Q

Suggest one reason why many waste disposal centres contain a separate section for cells and batteries

A

Prevent pollution of dangerous substances

51
Q

You are provided with the Daniel cell including a zinc electrode of a known mass
Briefly outline how you would carry out an experiment to confirm your answer to part b

A

Weigh mass of Zinc before and after
Allow cell to discharge until [Cu2+]—> 0.5
Confirm using colorimeter

52
Q

Suggest the main advantage of using hydrogen in a fuel cell rather than in an internal combustion engine

A

Less energy wasted

53
Q

Suggest two reasons why the concentratation of chlorine in the vase decreases with time
Chlorine used to kill bacteria

A

1- reacts with water to from HCl and HClO
2- used up killing bacteria

54
Q

Give two reasons why the value of the cell emf that you calculated in part b would not normally be achieved when the cell operates in sea water

A

Non-standard conditions and aluminium may react in an alkaline solution

55
Q

Explain briefly the results obtained for the iron foil (constant mass)

A

It is protected by the operation of a sacrificial anode

56
Q

Suggest briefly why the aluminium foil stopped decreasing in mass over time

A

Coating developed which prevents further erosion

57
Q

Explain the function of the salt bridge (2marks)

A

Allow the transfer of ions between electrodes
Maintain charge

58
Q

Suggest the major advantage of using the fuel cell compared to internal combustion engine

A

Internal combustion wastes more heat energy

Chemistry (25 decks)

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