3.1.11 electrode potentials and electrochemical cells

Question 1

what type of reaction takes place in electrochemical cells

Answer 1

redox reactions, where therefore e- are transferred from the reducing agent to the oxidising agent indirectly via an external circuit (wire)

Question 2

potential difference

Answer 2

difference in electrical potential between two electrodes. representing the energy required to move a unit charge, the PD is caused by the ability of electrons to flow from one half cell to the other (so the difference in reductive and oxidising ability allows the electric current to do work)

measured in volts

Question 3

what is a half cell/electrode

how do you represent a half cell

Answer 3

a basic half cell is a solid metal electrode immersed in a solution of its ions; eg iron metal in FeSO4 solution.

represented with a redox equilbrium. this is because the half cell on its own has an equilibrium. ALWAYS WRITTEN IN DIRECTION OF REDUCTION. so e- on left side of equation and oxidised form is getting reduced

Question 4

describe the equilibrium occurring at metal half cells

Answer 4

• metal atoms lose e- to become metal cations. they get oxidised
• and metal cations can lose e- to become metal atoms. they get reduced
• when eqm is reached the rate of metal atoms dissolving into metal ions is equal to the rate of ions become atoms. therefore the -ve charge on the metal (due to e- lost and left there) and the +ve charge of solution (cations) is constant
• e- are lost (forming oxidised form: ions) and picked up (forming reduced form, ions become atoms and deposit onto metal) at the solid metal electrode

Question 5

what is the electrode potential

Answer 5

indicates the tendency of a substance to be oxidised or reduced at an electrode.

it is the difference in charge between the -vely charged solid metal and the +vely charged (bc of metal ions) solution this charge separation is also the electrochemical potential. these charges are constant at eqm so the charge separation is constant

this can only be measured by the voltage between two cells. so it is compared to a reference electrode, the SHE

Question 6

EMF

voltage

potential difference

Answer 6

V reading under standard conditions

potential difference between two electrodes

aka V; represents energy required to move a unit charge. measured in V

Question 7

about the SHE’s electrode potential

why can it be used as a reference

Answer 7

the standard hydrogen electrode is the reference electrode used to determine indirectly the electrode potential of a half cell. we take the electrode potential of the SHE to be 0V, it is an arbitrary reference point. by definition, not necessarily reality, the electrode potential of the SHE is 0

because of this, we can form cells by attaching another electrode to the SHE and measuring the voltage reading. this is the E potential of the other electrode. all E potential values are relative to the SHE, therefore consistent and so comparable

Question 8

about the SHE’s set up

Answer 8

• Pt metal electrode extending from a wire
• H2 gas being bubbled into
• solution of HCl
• always left electrode (in conventional position of __ode)
• standard hydrogen electrode therefore conditions:
  1. standard pressure of H2 gas: 100kPa
  2. [H+] = 1moldm3
  3. 298K

Question 9

the same standard conditions must be applied to all electrodes having their standard electrode potential measured.

again, what are they, and what to look out for?

Answer 9

100kPa
1M solutions
298K
—–
the solutions must be 1moldm3 of the ION WE NEED. so if the iron we want is Fe3+ and the given sat is Fe2(SO4)3 (Fe III sulphate) the conc of the salt would be 0.5M because it released two moles Fe3+ for every mole salt

Question 10

whats the significance of the SHE always being on the left

Answer 10

it allows you to know about a meta electrodes oxidising/reducing tendencies from the sign

because if its positive than the more reactive electrode is on the left, as it should be. so the SHE is more reactive, so losing e-, so getting oxidised and e- are flowing from more negative electrode potential (of 0) to less negative; or more reactive to less reactive

and if its negative THIS HAS NOTHING TO DO WITH THE MAGNITUDE OF THE VOLTAGE. THIS DOES NOT MEAN OH ITS LESS THAN 0 ITS REALLY BAD. the number itself of the voltage is the magnitude of the voltage. the sign indicates direction of e- flow. if the sign is negative it means e- are flowing from, instead of as we would conventionally set it up, the conventional position of __ to that of __. so opposite to how we’d want it. indicating that the other electrode is more reactive than SHE, therefore e- are flowing FROM it TO the SHE; SHE gets reduced

Question 11

why are the conditions so important when measuring the electrode potential

Answer 11

the potential of a half cell is affect by changes in conc, temp an pressure. all of these aspects affect the ror and therefore the eqm. position of eqm changes if conditions change; and eqm determines the charge separation and therefore the electrode potential

electrode potentials need to be comparable so to make it fair they must all be measured under standard conditions

Question 12

how would a change in concentration affect an electrode/ cell

Answer 12

shifts to oppose the higher conc. [cation] is always 1moldm^3. if more than that, eqm will shift to decrease [cation] so shifts in direction of reduction (to metal atom). so electrode potential is more positive because less oxidation occuring

if less than that, eqm will shift to increase [cation] so shifts in direction of oxidation so more ngative E because more oxidation occuring

[[eqm positions shift to meet Kc, or more accurately, equilibria at certain temps fall into certain position which is described by Kc and therefore sysytem “adheres” to Kc but really its the opposite way]]

Question 13

[[a bit about kc thats unnecessary but cute]]

Answer 13

all equilibria at a certain temp are basc governed by their Kc, ofc not because they know to work in accordance with that number but by nature they settle at a certain eqilibrium and Kc describes this natural phenomenon mathematically

It doesn’t control the system — it just mathematically represents the natural balance the system settles into.

when a reversible reaction happens, the rates of the forward and backward reactions naturally adjust as concentrations change.

Eventually, the system reaches a steady state where the forward and backward rates are equal — this is equilibrium.

At that point, the concentrations of the substances fit the ratio defined by Kc.

Question 14

how would a change in temperature affect an electrode/ cell

Answer 14

ELECTRODE
depending on which way (red or ox) is exo or endo, will shift to oppose change in temp

CELL
depends on how each electrode responds to temp and how that might change the e transfer

Question 15

how would a change in pressure affect an electrode/ cell

Answer 15

of gas electrodes eg SHE
- h2 formed in red reaction
- so when pressure increased shifts to decrease, so more oxidation, so E becomes more -ve. so no longer 0 by definition. so can’t be standard (why the conditions so important)
- when decreased, shifts to increase so E becomes more positive as moves in direction of reduction.

Question 16

whats the point of the Pt electrode

characteristics that make it suitable

Answer 16

• inert therefore unreactive and won’t react with any of the species which would change its conc and affect the electrode potential, changing the conditions
• conducts electricity therefore electrons can flow in and out of it, allows flow of electrons and the redox equilibria in the SHE
• acts as an external electrical contact, electron sink and reserve, solid metal allowing e- to be transferred between H+ and H2
• often coated with fine Pt powder, increasing surface area so increased rate of reaction

Question 17

what happens at the Pt electrode

Answer 17

2H+ + 2e- <=> H2

H+ get reduced at the Pt electrode, gaining e- from it and becoming H2
H2 get oxidised at the Pt electrode, losing e- to it and becoming 2H+

Question 18

standard electrode potential definition

Answer 18

potential of an electrode measured against the SHE under standard conditions, that is to say: 1M for solutions, 298K and 100kPa

the difference in potential (V reading) between a given electrode and the SHE. measured in V

Question 19

set up of a cell

Answer 19

each electrode hooked up with leads and crocodile clips to a high resistance voltmeter

more negative E should go on the left

two solutions connected via a salt bridge. this is a piece of filter paper moistened with/ a glass U tube filled with saturated salt solution

Question 20

point of the high resistance voltmeter

Answer 20

• stops current flowing in the circuit
• current is the rate of flow of charge (negative charge of e- in an electrical circuit)
• as current flows the concs of ions begin to change, so no longer measuring E, which is what we wanted
• so you use this to stop current as much as pissible, so V reading is as representative of E as possible
• so you can measure PD which is what you want when determining standard electrode potentials.

Question 21

point of the salt bridge

Answer 21

• provides an ionic connection between 2 ionic solutions. ions are free to move in the salt bridge. normally made of KNO3.
• the anion (-ve, NO3-) is attracted to +ve charge. so they move to direction of electrode being oxidised, because e- going through wire and away. so more positive.
• the cation (+ve, K+) goes to electrode being reduced because losing +ve charge because metal cations in solution being reduced.
• this movement keeps each electrode electrically neutral so reaction (flow of e-) can continue
• completes circuit by preventing a buildup of charge. allowing continuous flow of e-

Question 22

cell notation - shorthand notatio to represent the redox reactions in an electrical cell

Answer 22

II - salt bridge
I - phase boundary, separates species in different phases
, - separates different species in same phase eg aq

• must be a solid at either end of the cell notation, whether this is Pt or another external electrical contact or the reduced form of the metal in a metal electrode
• if know more electrode with more -ve E, they go on left (so what would be the anode, gets oxidised)
• species should be arranged by ROOR; reduced forms further from salt bridge. oxidised forms nearer to the salt bridge
• include H+ AND H2O
• state symbols

Question 23

what happens to the redox equilibria when a cell is formed from two half cells/electrodes

Answer 23

equilibria collapse
it shifts very far to left or right (so in direction of oxidation (so being oxidised) or reduction (so being reduced))

this causes flow of e, this is caused by difference in potential

so extent of e- flow is then indicative of PD

Question 24

how does metal reactivity affect: the redox equilibria at its half cell;

the electrode potential of its half cell

Answer 24

• a more reactive metal loses e- more easily, getting oxidised to form ions. so forms ions more readily. so its eqm position is more equal or to the left because of this
• a less reactive metal forms ions less readily and for those that do, they are more likely to gain back electrons and become metal atoms. so its eqm position lies more to the right
• this means the more reactive a metal is, the more likely it will be oxidised and lose electrons. so its standard E potential will be more negative because that means it is more reactive than the SHE and the SHE gets reduced

Question 25

different types of electrodes

Answer 25

metal gas redox

Question 26

metal electrode

Answer 26

solid metal electrode in solution of its ions

Question 27

gas electrode

Answer 27

gas bubbled into solution of its ions no solid conductive surface so need inert metal electrode eg Pt electrode so e- can flow SHE is a gas electrode

Question 28

redox electrode

Answer 28

2 different ions of the same element. because charge is oxidation state, therefore one is the reduced form and one is the oxidised form eg Fe3+ and Fe2+ - they are in the same solution. - ion solutions must be 1M. look out for salts that when dissolved will release more than 1 ion - needs Pt electrode

Question 29

what does the size of the V tell you

Answer 29

BIG - this is like more than 1 - greater difference in potential, further apart in electrochemical series, Ecell is larger SMALL - like 0.1 or less - similar in reactivity, little difference in potential - weaker redox reaction so small flow of e- - near to each other in electrochemical series

Question 30

positive electrode meaning

Question 31

negative electrode meaning

Question 32

anode meaning

Question 33

cathode meaning

Question 34

what does the sign of the V tell you

Answer 34

has nothing to do with the magnitude. magntiude of V is the number itself. this tells you about the direction of the flow of e- 1. if +ve, cell is set up correctly. more negative electrode, so the anode, so the ___ electrode ,,, is connected to black lead. 2. if -ve, the more negative electrode is connected to red lead. so in place of where cathode (reduced, ___electrode) should be

Question 35

observations when have formed a cell

Answer 35

1. value on voltmeter 2. more -ve electrode wears away. because metal atoms are oxidised to ions so solid has reduced mass and size 3. less -ve electrode gets larger. because its metal cations in solution are reduced to metal atoms and deposit onto solid metal electrode. increase in size and mass 4. [ions in solution] decreases in less -ve electrode. might be able to tell by colour 5. [ions in solution] increase in more -ve electrode. might be able to tell by colour

Question 36

for any cell the EMF/V will be equal to...

Answer 36

the standard electrode potential of that cell

Question 37

how to calculate EMF

Answer 37

_E_cell = _E_right - _E_left _E_cell = _E_red - _E_ox _E_cell = _E_cathode - _E_anode

Question 38

when calculating EMF with eq *_E_cell = _E_right - _E_left * whats the assumption how can we account for this

Answer 38

that the left half cell is more -ve than the right (in E) so basc assuming that cell has been set up correctly ----- if we dont know what should go on right or left, we can use any of the following - anticlockwise rule - number line

Question 39

anticlockwise rule

Question 40

number line

Question 41

when calculating EMF with eq _E_cell = _E_red - _E_ox how can we know who's who

Answer 41

applying the redox theory we know about electrodes we know e- flow from electrodes with more negative E to electrodes with less negative so that means whoever has more negative E is getting oxidised, losing e- ,, which flow to other electrode

Question 42

what is the electrochemical series what can it be used for

Answer 42

a series of electrode potentials from +ve to -ve (can be either orientation) akin to the reactivity series of metals. all relative to the SHE, which is at 0 in the middle somewhere - identifying electrodes that would make good cells - determining feasibility of reactions ; predicting what reactions could occur - explaining why reactions happen

Question 43

what does a more positive E tell you on the electrochemical series; what does being on the more +ve end show

Answer 43

- increasing tendency to get reduced. they gain e- more easily. the metal ions get reduced to atoms - increasing tendency therefore to act as oxidising agents - so good/ better oxidising agents

Question 44

what does a more negative E tell you on the electrochemical series; what does being on the more -ve end show

Answer 44

- increasing tendency to get oxidised. they lose e- more easily. the metal atoms get oxidised to ions - increasing tendency therefore to act as reducing agents - so good/ better reducing agents

Question 45

how can we talk about the flow of e- in a cell? they flow from __ to __ [6]

Answer 45

e- flow from 1. the more -ve E to the less -ve 2. the anode to the cathode 3. the electrode undergoing ox to the one undergoing reduction 4. the more reactive metal/electrode to the less 5. the better reducing agent to the better oxidising agent 6. the __ electrode to the __ electrode

Question 46

how can the electrochemical series be used to identify components of the most powerful cell

Answer 46

powerful cells with have large EMFs EMF = _E_ cell the larger the difference between electrode potentials, the larger the EMF therefore can also be said that the better oxidising agent paired with the better reducing agent, the larger the EMF

Question 47

how can the electrochemical series be used to predict feasibility of reactions

Answer 47

theyre basc asking you will a reaction take place between two things on the electrochemical series. you can use the potentials to predict what redox reaction will take place between two species. unless sepcified, this doesnt have to be in context of cells. so we need to see what direction the eqms would go in (collapse into) when species put together. who gets oxidised and reduced then evaluate q stem

Question 48

how can the electrochemical series be used to explain why reactions happen

Answer 48

the reason a reaction between 2 electrodes takes place (why a voltage is measured) is because when connected, the better reducing agent has a higher tendency to lose e-; vv ,, and this causes a flow of e- and a redox reaction you can explain this in terms of the standard electrode potentials: x half cell (use electrode notation) has a more -ve _E_ than y half cell (^). so it is more reactive and gets oxidised more easily. it loses e-. [species in y half cell that is oxidised form] gains e- (becoming reduced) from [reduced form in x half cell] as it is oxidised.

Question 49

electrode notation

Answer 49

Fe(s)/Fe2+(aq) Fe3+/Fe2+ ^^exs

Question 50

secondary standard electrode meaning and types

Answer 50

The standard hydrogen electrode is difficult to use, so often a different standard is used which is easier to use. These other standards are themselves calibrated against the SHE. This is known as using a secondary standard- i.e. a standard electrode that has been calibrated against the primary standard. silver / silver chloride calomel electrode

Question 51

about the silver silver chloride 2ndary standard electrode

Answer 51

- calibrated against the she - safer because h2 is flammable and mercury causes brain damage - easier to make and sustain - made of Ag wire as electrode, coated with solid AgCl - this is submerged in solution of saturated KCl AgCl(s)+e−⇌Ag(s)+Cl−(aq) Ag(s)|AgCl(s),KCl(aq)

Question 52

about the calomel electrode

Answer 52

- calibrated against the she - safer because h2 is flammable ⚗️ What’s in it: - Mercury (Hg(l)) in contact with solid calomel (Hg₂Cl₂). - Dipped in a saturated KCl solution. ⚖️ The equilibrium: Hg₂Cl₂(s) + 2e⁻ ⇌ 2Hg(l) + 2Cl⁻(aq) The liquid mercury acts as the electron conductor. The solid Hg₂Cl₂ acts as a source/sink of Cl⁻ depending on redox direction. Hg(l)|Hg2Cl2(s),KCl(aq, sat'd )∥

Question 53

if they give you a cell setup and ask you to figure out Ecell if they give you half eqs and ask you to figure out Ecell

Answer 53

Ecell = Erhs - Elhs exactly as theyve given it ----- Ecell = Ered - Eox - do acw rule to see whos getting reduced/oxidised - or recall that more negative E = gets oxidised

Question 54

Electrochemical cells can be used...

Answer 54

Electrochemical cells can be used as a commercial source of electrical energy. this is the commercial applications of electrochemical cells.

Question 55

types of cell

Answer 55

rechargeable non rechargeable fuel

Question 56

what is a rechargeable cell

Answer 56

one in which the reactions occuring within can be reversed when an external current is applied. chemicals are regenerated types include lithium ion lead acid nickel cadmium

Question 57

lithium ion cell

Answer 57

discharging positive electrode: Li (in graphite)→Li+ + e- and negative electrode: Li+ + CoO2 + e– → Li+[CoO2]- when recharging these reactions are reversed

Question 58

lead acid batteries

Answer 58

- 6 cells, used in cars. on a larger scale cells provide energy to power a vehicle, as shown with these.

Question 59

nickel cadmium

Question 60

non rechargeable cells

Answer 60

- irreversible - one use - emf gets smaller until cell goes flat, emf=0 zinc carbon alkaline

Question 61

alkaline v zinc carbon nr cell

Answer 61

zc cheaper, standard one used alkaline more expensive but has longer life

Question 62

ZC nr cell might be wrogn

Answer 62

2MnO2(s)+H2O(l)+2e−→Mn2O2(s)+2OH−(aq) Zn(s)→Zn2+(aq)+2e−

Question 63

alkaline nr cell

Question 64

fuel cells

Answer 64

- continuous supply fo the chemicals into cell so none runs out - need a costant supply - as long as have that wont die or need recharging - most common is oxygen hydroegen fuel cell

Question 65

what occurs in a fuel cell

Answer 65

at anode, negative electrode: - h2 comes in - it is oxidised - the h+ ions move through electrolyte to cathode (positive) at cathode, positive electrode: - h+ reacts with o2 to form h2o in a reduction reaction electrons move between electrodes through a wire h2 is oxidised to h+ and o2 is reduced to h2o

Question 66

in alkaline hydrogen fuel cell

Question 67

in acidic hydrogen fuel cell

Question 68

waste issues assoicated with cells

Answer 68

nr cell: get eroded away the capsule so toxic chemicals get into soil and water r cells: ditto but less so because reusable fuel cells: hydrogen is gotten by fossil fuels so pollution

Question 69

benefits associated with cells

Answer 69

NR cheap R reusable so cheaper in long run less waste lower envtal impact FUEL CELLS water as waste dont need recharging v efficient

Question 70

badness of fuel cells

Answer 70

- needs constant supply - h2 flammable - h2 gotten by fossil fuels - high cost

Question 71

why would the cell emf of a nr cell fall over time and eventually reach 0

Answer 71

falls as reactants get used up so less oxidation and reduction going on reaches 0 because goes flat when one or boh reactants used up completely

Question 72

whats the main adv of a fuel cell over n/r cell (2)

Answer 72

doesnt go flat doesnt need recharging