electrode potentials Flashcards
def: electrode potential
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
def: half cells/ electrodes
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
if there is a large voltage, equilibrium is to the…
right
if there is a small voltage, equilibrium is to the…
left
why is a salt bridge (KNO3) used between the two beakers of electrolytes?
to complete the circuit, so the ions are free to move
unreactive with the electrodes
why might the current produced by a cell fall to serious after some time?
all reactants are used up
what happened once reactants are all used up?
stops working or stops to leak
left electrode:
negative electrode
oxidation always occurs at the left electrode
right electrode:
positve electrode
reductions always occur at the right electrode
electrons always flow from
negative (left) electrode to the
positive (right) electrode
platinum electrodes:
used when no other solid metals in reactions,
unreactive metals to complete the circuit,
conducts electricity
standard hydrogen electrode:
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
standard electrode potential:
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
conventional cell representation:
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
Equation for E cell:
E cell =
right - left
reduction - oxidation
more positive - least positive
def: cell discharge
when the E cell value is positive,
this means the reaction is feasible and the cell discharges -
produces a current
def: cell recharge
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
environmental advantage of rechargeable cells:
metals are reused
environmental disadvantage of rechargeable cells:
mains electricity is used to recharge, which may come from combusting fossil fuels, which releases CO2
why are conditions so important?
the equilibrium could shift to right:
more M and less M2+ and electrons,
fewer electrons the Eo would become more positive
why are conditions so important?
the equilibrium could shift to left:
less M and more M2+ and electrons,
more electrons the Eo would become more negative
rechargeable batteries - lithium ion battery
battery is discharged:
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
rechargeable batteries - lithium ion battery
battery is recharged:
Li + e- -> Li
reduction, right electrode
Li(CoO2) -> CoO2 + Li + e-
oxidation, left electrode
overall equation:
Li(CoO2) -> CoO2 + Li
def: fuel cells
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.
advantage of using fuel cells for energy rather than fossil fuels:
greater efficiency than burning hydrogen in a combustion engine
less-polluting as water is the only product
disadvantage of using fuel cells for energy rather than fossil fuels:
H2 is difficult to store
fossil fuels are combusted to produce the hydrogen, which releases carbon dioxide
advantages of fuel cells compared to other types or cells
voltage is constant, as fuel and oxygen is supplied constantly so concentrations of reactants remain constant.
def: electrolyte
can be acid or alkaline membrane which allows ions to move from one compartment of the cell to another,
acts as a salt bridge
hydrogen-oxygen fuel cell:
in the alkaline solution
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)
hydrogen-oxygen fuel cell:
in the acidic solution
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
