electrode potentials

Question 1

def: electrode potential

Answer 1

2 different metals places in a salt solution and join them together, electrons will flow from more reactive metal to the less reactive metal, current produced can be used to power electrical devices

Question 2

def: half cells/ electrodes

Answer 2

a strip of metal dipped into a solution of its own ions

• equilibrium is set up:
  equil takes place on surface of solid metal, at the interface between the two phases:
  Cu2+ ions dissolve in sol, give sol positive charge,
  electrons collect into the copper strip, gives strip negative charge,
  potential difference is established between the two

Question 3

if there is a large voltage, equilibrium is to the…

Answer 3

right

Question 4

if there is a small voltage, equilibrium is to the…

Answer 4

left

Question 5

why is a salt bridge (KNO3) used between the two beakers of electrolytes?

Answer 5

to complete the circuit, so the ions are free to move
unreactive with the electrodes

Question 6

why might the current produced by a cell fall to serious after some time?

Answer 6

all reactants are used up

Question 7

what happened once reactants are all used up?

Answer 7

stops working or stops to leak

Question 8

left electrode:

Answer 8

negative electrode
oxidation always occurs at the left electrode

Question 9

right electrode:

Answer 9

positve electrode
reductions always occur at the right electrode

Question 10

electrons always flow from

Answer 10

negative (left) electrode to the
positive (right) electrode

Question 11

platinum electrodes:

Answer 11

used when no other solid metals in reactions,
unreactive metals to complete the circuit,
conducts electricity

Question 12

standard hydrogen electrode:

Answer 12

voltage of standard hydrogen half cell is zero;
298K
H2 gas pumped in at 100kPa
electrolyte contains H+ ions of concentration 1 moldm-3
platinum electrode

Question 13

standard electrode potential:

Answer 13

list of electrochemical electrode equations - all shown as reduction
all species on left of arrow are oxidising agents;
all species on the right of arrow are reducing agents;
SOWR - SO has most positive Eo value

Question 14

conventional cell representation:

Answer 14

phase boundary
|| salt bridge
right electrode, more positive standard electrode potential then left electrode;
species with highest oxidation state should be written closest to salt bridge;
R|O||O|R

Question 15

Equation for E cell:

Answer 15

E cell =
right - left
reduction - oxidation
more positive - least positive

Question 16

def: cell discharge

Answer 16

when the E cell value is positive,
this means the reaction is feasible and the cell discharges -
produces a current

Question 17

def: cell recharge

Answer 17

if the reaction is reversible, the cell can be recharged by plugging it into the mains,
reverse reaction with occurs when the cell is recharged

Question 18

environmental advantage of rechargeable cells:

Answer 18

metals are reused

Question 19

environmental disadvantage of rechargeable cells:

Answer 19

mains electricity is used to recharge, which may come from combusting fossil fuels, which releases CO2

Question 20

why are conditions so important?
the equilibrium could shift to right:

Answer 20

more M and less M2+ and electrons,
fewer electrons the Eo would become more positive

Question 21

why are conditions so important?
the equilibrium could shift to left:

Answer 21

less M and more M2+ and electrons,
more electrons the Eo would become more negative

Question 22

rechargeable batteries - lithium ion battery
battery is discharged:

Answer 22

Li -> Li+ + e-
Li has been oxidised = neg, left electrode

CoO2 + Li+ + e- -> Li(CoO2)
Co has been reduced = pos, right electrode

overall equation:
CoO2 + Li -> Li(CoO2)
conventional cell representation:
Li|Li+||Li+, CoO2|Pt

Question 23

rechargeable batteries - lithium ion battery
battery is recharged:

Answer 23

Li + e- -> Li
reduction, right electrode

Li(CoO2) -> CoO2 + Li + e-
oxidation, left electrode

overall equation:
Li(CoO2) -> CoO2 + Li

Question 24

def: fuel cells

Answer 24

uses energy from the reaction of a fuel with oxygen to create a voltage
- combusting a fuel can create a flow of electrons

both fuel and oxidant are used up during reaction and need to be continuously provided if the cell is to continue to provide a constant voltage

modern fuel cells are based on hydrogen, or hydrogen-rich fuels such as methanol, natural gas and petrol.

Question 25

advantage of using fuel cells for energy rather than fossil fuels:

Answer 25

greater efficiency than burning hydrogen in a combustion engine
less-polluting as water is the only product

Question 26

disadvantage of using fuel cells for energy rather than fossil fuels:

Answer 26

H2 is difficult to store
fossil fuels are combusted to produce the hydrogen, which releases carbon dioxide

Question 27

advantages of fuel cells compared to other types or cells

Answer 27

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

Question 28

def: electrolyte

Answer 28

can be acid or alkaline membrane which allows ions to move from one compartment of the cell to another,
acts as a salt bridge

Question 29

hydrogen-oxygen fuel cell:
in the alkaline solution

Answer 29

at anode, neg electrode:
H2(g) + 2OH-(aq) <—> 2H2O(l) + 2e-
= oxidation

at cathode, pos electrode:
1⁄2 O2(g) + H2O(l) + 2e- <—>2OH-(aq) = reduction

overall equation:
H2(g) + 1⁄2 O2(g) → H2O(l)

Question 30

hydrogen-oxygen fuel cell:
in the acidic solution

Answer 30

at anode, neg electrode:
H2(g) <—> 2H+(aq) + 2e-
= oxidation

at cathode, pos electrode:
1⁄2 O2(g) + 2H+(aq) + 2e- <—> H2O(l)

overall equation:
H2(g) + 1⁄2 O2(g) → H2O(l)
pure hydrogen only emits water