Physical Unit 1.11: Electrode Potentials & Cells (Redox Equilibria) Flashcards
describe the +ve & -ve potential of an electrode in terms of equilibria
electrode w -ve potential:
when equlibrium lies to the left, the metal becomes -vely charged due to e-s being released & building up on the metal
electrode w +ve potential:
when equilibrium lies to the right, the metal becomes +vely charged due to e-s being used up to form metal from metal ions
what does the position of equilibrium depend on?
the metal
reactive metals tend to form Mn+ ions so -ve charge builds up on the metals
more reactive metals have -ve potentials
more unreactive metals have +ve potentials
define electrode/half-cell
a metal dipping into a solution of its ions
what electrode is used when there is no solid metal involved in the half-equation & why?
pure platinum electrode
inert
a solid metal is needed to allow the flow of e-s
what are the 3 types of electrodes?
metal electrodes
gas electrodes
redox electrodes
describe metal electrodes
metal solution surrounded by a solution of its ions
describe gas electrodes
for a gas & a solution of its ions
inert metal (usually Pt) is the electrode to allow the flow of e-s
describe redox electrodes
for 2 different ions of the same element
2 types of ion are present in solution with an inert metal electrode (usually Pt) to allow the flow of e-s
e.g. Pt(s)|Fe2+(aq)|Fe3+(aq)
how do you measure the potential of an electrode?
connect half-cell to another half-cell of known potential & measure the potential difference b/w the 2 half-cells
the standard hydrogen electrode (SHE) is assigned the potential 0 & is the primary standard = the potential to which all others are compared/measured
so unknown half-cell connected to SHE with SHE on the left
what is an electrochemical cell?
2 half-cells connected
what does primary standard mean?
the potential to which all others are compared/measured
describe how 2 half-cells are joined together to give a complete circuit
2 metal electrodes are joined with a wire - to allow the transfer of e-s
the metals are dipped into a solution of their ions
the 2 solutions are joined with a salt bridge - to allow the flow of ions
a voltmeter (high resistance so low current) is often included to measure the potential difference (emf)
what can a salt bridge be made of?
a piece of filter paper soaked with a solution of unreactive ions
a tube containing unreactive ions in agar gel
e.g. KNO3 is often used in a salt bridge as K+ & NO3- are quite unreactive
what are the standard conditions?
1.0moldm-3 solution of the ions (NB for redox electrode, 1.0moldm-3 for each ion)
298K
100kPa if half-cell includes gases
pure metal electrode
why are standard conditions required?
the position of the redox equilibrium changes when conditions change
how will the emf change if [a condition that affect position of equilibrium] is increased/decreased?
either:
emf increases/more +ve
as equilibrium shifts right so e-s used up
so more reduction happens
or
emf decreases/more -ve
as equilibrium shifts left so e-s released
so more oxidation happens
draw standard hydrogen electrode
what is the half equation?
what is the cell notation?
see booklet
2H+(aq) + 2e- <–> H2(g) convention to write as reduction
Pt(s)|H2(g)|H+(aq)
how is the conventional representation of cells written?
RO||OR
oxidation||reduction
|| indicates salt bridge
| indicates phase boundary - if no phase change, use a comma
solid metal electrode on the outside
emf =
Eθright - Eθleft
when measuring Eθ vs SHE, SHE always on the left
define secondary standard & why are they used?
= another standard electrode that has been calibrated against the SHE
SHE is difficult to use bc uses a gas, H2 is flammable & hard to keep at 100kPa
describe the redox process
metal atoms lose e-s at the -ve electrode/anode = oxidation
these e-s travel through the wire to the +ve electrode/cathode & are gained by ions to produce metal atoms = reduction
+ve electrode normally RHS
emf is not affected by the # of e-s in equation
what must be true for a reaction to be feasible?
emf must be +ve
Eθ ox. agent > Eθ red. agent
[…] is a better oxidising agent thatn […]
if emf is -ve?
feasible in reverse direction
what happens when a cell is discharged vs recharged?
discharged: feasible reaction occurs, emf is +ve
recharged: opposite, non-feasible reaction occurs, emf is -ve
why does the emf of a cell change when electrodes are connected & current flows?
conditions are no longer standard
the conc. of ions changes
what is the electrochemical series?
describe
list of electrode potentials in order of decreasing or increasing potential
written as reductions
+ve potentials = good at attracting e-s so good oxidising agents
-ve potentials = good at releasing e-s so good reducing agents
flip & add
flip more -ve potential if have the choice
can […] oxidise […]?
emf must be +ve
when to use flip & add vs when to use R-L to calculate emf
if given cell notation, do R-L
if given equations, do flip & add
what are the different types of cell?
non-rechargeable
rechargeable
fuel cells
define battery
more than one cell joined together
describe non-rechargeable cells
chemicals are used up over time & emf decreases
when 1 or more of the chemicals have been used up, the cell is flat & emf = 0
cells cannot be recharged so are single use & must be disposed off
(graph of emf vs time for non-rechargeable cells)
see booklet
describe rechargeable cells
reactions are reversible - reversed by applying an external current & regenerate the chemicals
what are the equations for lithium ion (rechargeable) cell?
used in phones, laptops etc.
Li+ + CoO2 + e- <–> LiCoO2 +0.60V
Li+ + e- <–> Li -3.00V
emf = +3.60V
overall reaction during discharge: Li + CoO2 –> LiCoO2
Eθ is +ve
overall reaction during recharge:
LiCoO2 –> Li + CoO2
Eθ is -ve - not feasible so more than 3.60V needs to be applied
(graph of emf vs time for rechargeable cells)
see booklet
describe fuel cells
have continuous supply of chemicals into the cell so do not run out of chemicals & do not need recharging
what is the most common fuel cell?
hydrogen-oxygen fuel cell
diagram for hydrogen-oxygen fuel cell
see booklet
what is the overall equation for hydrogen-oxygen fuel cells?
2H2 + O2 –> 2H2O
emf +1.23V
same equation & emf in alkaline & acidic conditions
what are the equations for the reactions at the -ve & +ve electrodes of hydrogen-oxygen fuel cells in alkaline conditions?
at -ve electrode: ox.
H2 + 2OH- –> 2H2O + 2e-
at +ve electrode: red.
O2 + 2H2O + 4e- –> 4OH-
what are the equations for the reactions at the -ve & +ve electrodes of hydrogen-oxygen fuel cells in acidic conditions?
at -ve electrode: ox.
H2 –> 2H+ + 2e-
at +ve electrode: red.
O2 + 4H+ + 4e- –> 2H2O
what is the cell notation for the hydrogen-oxygen fuel cell in alkaline conditions?
Pt(s)|H2(g)|OH-(aq),H2O(l)||O2(g)|H2O(l),OH- (aq)|Pt(s)
what is the cell notation for the hydrogen-oxygen fuel cell in acidic conditions?
Pt(s)|H2(g)|H+(aq)||O2(g)|H+(aq),H2O(l)|Pt(s)
are hydrogen-oxygen fuel cells efficient & what is the waste product?
yes
water is the only waste product
what is the formula for efficiency?
useful energy output/total energy input
table of benefits & risks of using different types of cells
see booklet
