Physical Unit 1.11: Electrode Potentials & Cells (Redox Equilibria)

Question 1

describe the +ve & -ve potential of an electrode in terms of equilibria

Answer 1

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

Question 2

what does the position of equilibrium depend on?

Answer 2

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

Question 3

define electrode/half-cell

Answer 3

a metal dipping into a solution of its ions

Question 4

what electrode is used when there is no solid metal involved in the half-equation & why?

Answer 4

pure platinum electrode
inert
a solid metal is needed to allow the flow of e-s

Question 5

what are the 3 types of electrodes?

Answer 5

metal electrodes
gas electrodes
redox electrodes

Question 6

describe metal electrodes

Answer 6

metal solution surrounded by a solution of its ions

Question 7

describe gas electrodes

Answer 7

for a gas & a solution of its ions
inert metal (usually Pt) is the electrode to allow the flow of e-s

Question 8

describe redox electrodes

Answer 8

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)
standard conditions: both ion solutions must be 1 mol dm-3, so [Fe2+] = 1 moldm-3 & [Fe3+] = 1 moldm-3

Question 9

how do you measure the potential of an electrode?

Answer 9

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

Question 10

what is an electrochemical cell?

Answer 10

2 half-cells connected

Question 11

what does primary standard mean?

Answer 11

the potential to which all others are compared/measured

Question 12

describe how 2 half-cells are joined together to give a complete circuit

Answer 12

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)

Question 13

what can a salt bridge be made of?

Answer 13

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

Question 14

(why are K+ ions unreactive?)

Answer 14

K+ unreactive because e- configuration is …3p6 full outer shell of e-s, which is a v stable, low-energy configuration
K+ fully oxidised

Question 15

what are the standard conditions?

Answer 15

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

Question 16

why are standard conditions required?

Answer 16

the position of the redox equilibrium changes when conditions change

Question 17

how will the emf change if [a condition that affect position of equilibrium] is increased/decreased?

Answer 17

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

Question 18

draw standard hydrogen electrode
what is the half equation?
what is the cell notation?

Answer 18

see booklet
2H+(aq) + 2e- <–> H2(g) convention to write as reduction
Pt(s)|H2(g)|H+(aq)

Question 19

how is the conventional representation of cells written?

Answer 19

RO||OR
oxidation||reduction
|| indicates salt bridge
| indicates phase boundary - if no phase change, use a comma
solid metal electrode on the outside

Question 20

emf =

Answer 20

Eθright - Eθleft
when measuring Eθ vs SHE, SHE always on the left

Question 21

define secondary standard & why are they used?

Answer 21

= 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

Question 22

describe the redox process

Answer 22

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

Question 23

emf is not affected by the # of e-s in equation

Question 24

what must be true for a reaction to be feasible?

Answer 24

emf must be +ve
Eθ ox. agent > Eθ red. agent
[…] is a better oxidising agent than […]

Question 25

if emf is -ve?

Answer 25

feasible in reverse direction

Question 26

what happens when a cell is discharged vs recharged?

Answer 26

discharged: feasible reaction occurs, emf is +ve recharged: opposite, non-feasible reaction occurs, emf is -ve

Question 27

why does the emf of a cell change when electrodes are connected & current flows?

Answer 27

conditions are no longer standard the conc. of ions changes

Question 28

what is the electrochemical series? describe

Answer 28

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

Question 29

flip & add

Answer 29

flip more -ve potential if have the choice

Question 30

can [...] oxidise [...]?

Answer 30

emf must be +ve

Question 31

when to use flip & add vs when to use R-L to calculate emf

Answer 31

if given cell notation, do R-L if given equations, do flip & add

Question 32

what are the different types of cell?

Answer 32

non-rechargeable rechargeable fuel cells

Question 33

define battery

Answer 33

more than one cell joined together

Question 34

describe non-rechargeable cells

Answer 34

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

Question 35

(graph of emf vs time for non-rechargeable cells)

Answer 35

see booklet

Question 36

describe rechargeable cells

Answer 36

reactions are reversible - reversed by applying an external current & regenerate the chemicals

Question 37

what are the equations for lithium ion (rechargeable) cell?

Answer 37

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

Question 38

(graph of emf vs time for rechargeable cells)

Answer 38

see booklet

Question 39

describe fuel cells

Answer 39

have continuous supply of chemicals into the cell so do not run out of chemicals & do not need recharging

Question 40

what is the most common fuel cell?

Answer 40

hydrogen-oxygen fuel cell

Question 41

diagram for hydrogen-oxygen fuel cell

Answer 41

see booklet

Question 42

what is the overall equation & emf for hydrogen-oxygen fuel cells?

Answer 42

2H2 + O2 --> 2H2O emf +1.23V same equation & emf in alkaline & acidic conditions

Question 43

what are the equations for the reactions at the -ve & +ve electrodes of hydrogen-oxygen fuel cells in alkaline conditions?

Answer 43

at -ve electrode: ox. H2 + 2OH- --> 2H2O + 2e- at +ve electrode: red. O2 + 2H2O + 4e- --> 4OH-

Question 44

what are the equations for the reactions at the -ve & +ve electrodes of hydrogen-oxygen fuel cells in acidic conditions?

Answer 44

at -ve electrode: ox. H2 --> 2H+ + 2e- at +ve electrode: red. O2 + 4H+ + 4e- --> 2H2O

Question 45

what is the cell notation for the hydrogen-oxygen fuel cell in alkaline conditions?

Answer 45

Pt(s)|H2(g)|OH-(aq),H2O(l)||O2(g)|H2O(l),OH- (aq)|Pt(s)

Question 46

what is the cell notation for the hydrogen-oxygen fuel cell in acidic conditions?

Answer 46

Pt(s)|H2(g)|H+(aq)||O2(g)|H+(aq),H2O(l)|Pt(s)

Question 47

are hydrogen-oxygen fuel cells efficient & what is the waste product?

Answer 47

yes water is the only waste product

Question 48

what is the formula for efficiency?

Answer 48

useful energy output/total energy input

Question 49

table of benefits & risks of using different types of cells

Answer 49

see booklet

Question 50

what is the function of a salt bridge?

Answer 50

MS: the ions in the ionic substance in the salt bridge move through the salt bridge to complete the circuit

Question 51

why is ___ cell not recharged?

Answer 51

reaction cannot be reversed named product cannot be reduced to named reagent by applying an external current difficult to replace NAMED REAGENT