topic 14: redox II

Question 1

what does an electrochemical cell consist of

Answer 1

• a cell has two half–cells
• the two half cells have to be connected with a salt bridge
• simple half cells will consist of a metal (acts an
  electrode) and a solution of a compound containing
  that metal (eg Cu and CuSO4) - 1 moldm-3
• these two half cells will produce a small voltage if connected into a circuit. (i.e. become a battery or cell)
• standard conditions
• electrodes connected to a high resistance voltmeter
• LHS - oxidised
• RHS - reduced

Question 2

why does a voltage form

Answer 2

• when connected together the zinc half-cell has more of a tendency to
  oxidise to the Zn2+ ion and release electrons than the copper half-cell
  (Zn Zn2+ + 2e-)
• more electrons will therefore build up on the zinc electrode than the copper electrode
• a potential difference is created between the two electrodes. the zinc strip is the negative terminal and the copper strip is the
  positive terminal

Question 3

why use a high resistance voltmeter

Answer 3

• the voltmeter needs to be of very
  high resistance to stop the current
  from flowing in the circuit. in this
  state it is possible to measure the
  maximum possible potential difference (E).
• the reactions will not be occurring
  because the very high resistance
  voltmeter stops the current fromflowing

Question 4

what is the salt bridge used for

Answer 4

the salt bridge is used to connect up the circuit. the free moving ions conduct the charge

Question 5

what is a salt bridge made up of

Answer 5

• a salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate
  
  • the salt should be unreactive with the electrodes and electrode solutions
    ——-> e.g. potassium chloride would not
    be suitable for copper systems as chloride ions can form complexes with copper ions

Question 6

why is a wire not used in place of a salt bridge

Answer 6

because the metal wire would set up its own electrode system with the solutions and wires do not allow the flow of ions

Question 7

what would happen if the voltmeter is removed and replaced with a bulb

Answer 7

a current flows. the
reactions will then occur separately at each electrode. the voltage will fall to zero as the reactants are used up

Question 8

how can electrochemical cells be represented

Answer 8

cell diagram
Zn(s) | Zn2+ (aq) | | Cu2+ (aq) | Cu (s)

Question 9

what does each part of the cell diagram represent

Answer 9

• the solid vertical line represents the boundary between phases e.g. solid (electrode) and solution (electrolyte) * the double line represents the salt bridge between the two half cells
• the voltage produced is indicated
• the more positive half cell is written on the right if possible (but this is not essential)

Question 10

what electrode is used if the system doesn’t have an electrode and why

Answer 10

• platinum electrode
• provides a conducting surface for electron transfer
• a platinum electrode is used because it is unreactive and can conduct electricity

Question 11

how do you measure the electrode potential of a cell

Answer 11

• it is not possible to measure the absolute potential of a half electrode on its own. it is
  only possible to measure the potential difference between two electrodes
• to measure it, it has to be connected to another half-cell of known potential, and the potential difference between the two half-cells measured
• by convention we can assign a relative potential to each electrode by linking it to a reference electrode (hydrogen electrode), which is given a potential of zero volts

Question 12

what is the potential of all electrodes measured by comparing

Answer 12

comparing their potential to that of the standard hydrogen electrode

Question 13

what are the components of a standard hydrogen electrode (SHE)

Answer 13

1. hydrogen gas at pressure of 100 kPa
2. solution containing the hydrogen ion at 1.00 mol dm-3 (solution is usually 1 mol dm-3 HCl)
3. temperature at 298K

Question 14

describe how to draw an electrode set up

Answer 14

• on the right hand side, the electrode getting reduced is drawn in 1 mol dm-3 of a corresponding solution - anode - positive terminal
• on the left hand side, the electrode is getting oxidised in 1mol dm-3 of the corresponding solution - cathode - negative terminal
• salt bridge between each beaker
  high resistance voltmeter drawn
• the direction of current goes from LHS to RHS
  WRITE OUT THE STATNDARD CONDITIONS

Question 15

what is standard electrode potential

Answer 15

when an electrode system is connected to the hydrogen electrode system, and standard conditions apply the potential difference measured

Question 16

what is the equation for working out the EMF of a cell

Answer 16

E(rhs) - E(lhs)

Question 17

what does it mean for the more negative half cell

Answer 17

it is oxidised so it goes the opposite way

Question 18

what happens as the standard electrode potential increases and becomes more positive

Answer 18

tendency for
species on left to
reduce, and act
as oxidising
agents

Question 19

what happens as the standard electrode potential gets more negative

Answer 19

tendency for
species on right
to oxidise, and
act as reducing
agents

Question 20

what is Ecell directly proportional to

Answer 20

total entropy change and lnK
positive Ecell will lead to a positive total entropy change

Question 21

what is Ecell a measure of

Answer 21

a measure of how far from equilibrium the
cell reaction lies. the more positive the Ecell the more likely the reaction is to occur

Question 22

what is the effect of concentration on Ecell

Answer 22

as the concentration of the reactants increases so will the Ecell

Question 23

what is the effect of temperature on Ecell

Answer 23

most cells are exothermic in the spontaneous direction so applying Le Chatelier to a temperature rise to these would result in a decrease in Ecell because the equilibrium reactions would shift backwards

Question 24

what is a fuel cell

Answer 24

uses the energy from the reaction of
a fuel with oxygen to create a voltage

Question 25

what is the overall equation of a hydrogen fuel cell

Answer 25

2e- + 2H+ —> H2 E=0V
4e- + 4H+ +O2 —> 2H2O E=1.23V
in acidic conditions these are the
electrode potentials. the Ecell is the same as alkaline conditions as
the overall equation is the same
overall 2H2 + O2 —> 2H2O E=1.23V

Question 26

what are some advantages of hydrogen fuel cells

Answer 26

• less pollution and less CO2 (pure hydrogen emits only water whilst hydrogen-rich fuels
  produce only small amounts of air pollutants and
  CO2)
• greater efficiency
• hydrogen is readily available by the
  electrolysis of water

Question 27

what are some limitations of hydrogen fuel cells

Answer 27

• expensive
• storing and transporting hydrogen, in terms of safety, feasibility of a pressurised liquid and a limited
  life cycle of a solid ‘adsorber’ or ‘absorber’
• limited lifetime (requiring regular replacement and
  disposal) and high production costs
• use of toxic chemicals in their production

Question 28

what are the different ways that hydrogen can be stored in fuel cells

Answer 28

• as a liquid under pressure
• adsorbed on the surface of a solid material
• absorbed within a solid material

Question 29

describe advantages of ethanol/methanol fuel cells

Answer 29

ethanol fuel cells have also been developed. compared to hydrogen fuel cells they have certain advantages including…
- ethanol can be made from renewable sources in a carbon neutral way
- raw materials to produce ethanol by fermentation are
abundant
- ethanol is less explosive and easier to store than hydrogen
- new petrol stations would not be required as ethanol is a liquid fuel

Question 30

what are the equations that occur in an ethanol/methanol fuel cells

Answer 30

equation that occurs at oxygen electrode
4e- + 4H+ +O2 —> 2H2O E=1.23V
equation that occurs at ethanol electrode
C2H5OH + 3O2 → 2CO2 + 3H2O
overall equation
C2H5OH + 3H2O → 2CO2 + 12H+ + 12e-

Question 31

manganate redox titration

Answer 31

• the acid is needed to supply the 8H+ ions. some acids are not suitable as they set up alternative redox reactions and hence make the titration readings inaccurate.
  only use dilute sulfuric acid for manganate titration.
  insufficient volumes of sulfuric acid will mean the solution is not acidic enough and MnO2 will be produced instead of Mn2+
  MnO4-(aq) + 4H+(aq)+ 3e- —> MnO2 (s) + 2H2O
  the brown MnO2 will mask the colour change and lead to a greater (inaccurate) volume of manganate being
  used in the titration
• using a weak acid like ethanoic acid would have the same effect as it cannot supply the large amount of hydrogen ions needed (8H+)
• it cannot be conc HCl as the Clions would be oxidised to Cl2 by MnO4- as the Eo MnO4-/Mn2+ > Eo Cl2/Cl-
  MnO4-(aq) + 8H+(aq)+ 5e– —>Mn2(aq) + 4H2O(l) E+1.51V
  Cl2(aq) +2e– —> 2Cl–(aq) E +1.36V this would lead to a greater volume of manganate being used and poisonous Cl2 being produced
• it cannot be nitric acid as it is an oxidising agent. it oxidises Fe2+ to Fe3+ as Eo NO3-/HNO2> Eo Fe3+/Fe2+
  NO3- (aq) + 3H+(aq) + 2e– —> HNO2(aq) + H2O(l) Eo +0.94V
  Fe3+ (aq)+e–  Fe2+ (aq) Eo +0.77 V
  this would lead to a smaller volume of manganate being used

Question 32

in the redox titration between Fe2+ and MnO4- what is the colour change when the manganate is in the burette

Answer 32

colourless to pink

Question 33

what can cells be

Answer 33

non-rechargeable (irreversible), rechargeable and fuel cells