3.1.11 Electrode Potentials And Electrochemical Cells Flashcards
(30 cards)
Salt bridge
- normally filter paper saturated with sodium nitrate solution
- allows the flow of ions
- not made of wire - as would set up own electrode system
- why unreactive aq solution - so doesn’t react and mess experiments up
Voltmeter
Incredibly high resistance
To prevent the current from flowing in the circuit.
Voltage = electrodpotential
If want electrons to flow = replace with bulb
Reduction and oxidation in the cell (electron movement rule)
No voltmeter
Electrons more neg to more positive halfcell
Half cell
Metal
Dipped in its aq solution/ ions
Normally metal sulphate aq
NOPR rule
NO
PR
- more negative halfcell (more neg potentail difrence) = oxidised
- more positive halfcell = reduced
When half cell is reduced half equation
When reduce start with the ion !!!
Then go to the atom
(Oposite for oxidised)
Cell diagram
- 2 vertices line in middle = salt bridge
- Then next to the salt brifmdge - the most oxidised form (highest oxidation state
- Single vertices lone = change of state
Best practice
More positive half cell = RHS
More negative half cell =LHS
No solid conducting surface
Use a platinum electrode
- as need an electrode conducting surface
- platinum is unreactive
- conducts electricity
- changes the cell diagram
- do it normally but common between the two aq species and add py electrode on the end
Reasons for the standard hydrogen electrode
Can not measure the electrode potential of one half cell
Can only calculate the potential difference
Standard hydrogen electrode - refrence electrode -0V
Standard hydrogen electrode format
Platinum electrode
Solution needs proton often strong acid fully dissociated = eg HCl
H2 gas bubbled over the solution and electrode
Standard hydrogen electrode equation
At eqm
Can be oxidised or reduced depending on the other electrode attached
H2 -> <- 2H+ +2e-
Standard conditions for hydrogen electrode
H2 gas pressure 100kpa
Conc of H+ 1moldm3
Temp of 298K
Pt electrode
Conventional cell dugram for the standard hydrogen electrode
Pt(S)/ H2(g) /H+(aq) //
Standard hydrogen electrode trick to be carful of
When H2SO4 is the aq component
Diuretic base goes to 2mol H+
Meaning conc is NOT 1moldm-3
Needs to be 0.5 moldm-3
Standard conditions
All ions in solution 1 mol dm-3( if 2 aq component they are still both qmoldm-3)
If gas pressure of 100kpa
Temp 298K
No surrent flowing
Ecell =
= E reduced - E oxidised
= E right - E left
Strongest reducing / oxidising agent in electrochemical series
For finding oxidised
The most positive and the atom on the left
For finding reducing agent
The most negative half cell the atom
The atom on tye top right
Use electrode potential to explain why reaction takes place
Electrode potential of the(f2,f-) is greater than the electrode peotential (o2/h20)
this therfore f2 will oxidise the h20 to o2
Written as (f2/f-) as display halfcell = most oxidised form/ the next most oxidised form
Changing of conditions of the cell
Increase concentration of reactant = increase Ecell (increase reduced )
Decrease the conc of reactant = decrease Ecell
Most electrode cells in the spontaneouse direction are in teh exothermic direction
What is the electrochemical series
List of electrode potentials in numerical order
Suggest why two tm complexes of same metal have different electrode potentials
Different ligands
Why electrode potential of the standard hydrogen electrode is 0
By definition
Conventional representation of the alkaline hydrogen oxygen fuel cell
Pt/H2(g)/OH-(aq)/H2O(l)
//
O2(g)/H2O(l)OH-(aq)/Pt
Why the emf of a hydrogen oxygen file cell in acidic conditions is teh same as in alkaline
As the overall reaction is the same
2H2 + O2 -> 2H2O
