23. Redox and Electrode Potentials

Question 1

How do you balance half equations in basic solutions?

Answer 1

balance thing being oxidised/reduced

deduce e- needed from oxidation state change

balance on other side w OH-

balance on other side w H2O

Question 2

What type of reactions happen in batteries (voltaic cells)?

Answer 2

chemical reactions that transfer e- from one species to another

Question 3

When is a platinum electrode used?

Answer 3

when there is a mixture of 2 different ions of the same element or a gas and a solution of its ion

is used to transport the e-

Question 4

What does the standard electrode potential represent?

Answer 4

the tendency for the species in a half cell to either gain or lose electrons

Question 5

Define standard electrode potential.

Answer 5

the electromotive force of a half cell compared to the standard hydrogen electrode measured at 298K, solution conc. 1 moldm-3 and at a pressure of 100kPa

Question 6

Draw the standard hydrogen half cell connected to a Cu half cell.

Answer 6

Page 3 of electrode potentials leaflet

Question 7

What is a salt bridge and what is its purpose?

Answer 7

connects the 2 half cells

made of KNO3 as all compounds formed are soluble and it will not react w other solutions

allows ions to transfer between half cells to keep them neutral

Question 8

What component connects the 2 electrodes in each half cell other than the salt bridge?

Answer 8

high resistance voltmeter: measures potential for electron transfer

Question 9

In which direction are all half equations written when showing their E(theta)?

Answer 9

reduction half equations

Question 10

What does decreasing E(theta) mean?

Answer 10

the more negative E is, the greater its tendency to lose e- and go in the reverse direction (i.e. species on the right gets oxidised)

more negative = species on right is strong reducing agent

Question 11

What does increasing E(theta) mean?

Answer 11

the more positive E is, the greater its tendency to gain e- and go in the forward direction (i.e. as is written, species on left gets reduced)

more positive = species on the left is strong oxidising agent

Question 12

How do you calculate overall cell potential (E(cell))?

Answer 12

E(most +ve electrode) - E(most -ve electrode)

Question 13

If E of half cell 1 = 0.34V and E of half cell 2 = -0.11V which will be the forwards reaction?

Answer 13

half cell 1 forwards: species on left is strong oxidising agent
half cell 2 backwards: species on right is strong reducing agent

Question 14

At which electrode does oxidation take place?

Answer 14

anode: negative electrode: in half cell with most -ve E

Question 15

At which electrode does reduction take place?

Answer 15

cathode: positive electrode: in half cell with most +ve E

Question 16

At which electrode are electrons produced?

Answer 16

the electrode where oxidation is taking place: the negative electrode

Question 17

Which direction do electrons flow?

Answer 17

from the -ve (where they are produced) to the +ve electrode

Question 18

Zn2+(aq) + 2e- <=> Zn(s)
E(theta) = -0.76V
What will happen if the concentration of Zn2+ is increased?

Answer 18

equilibrium shifts right, removing e- from the system

electrode potential becomes less negative

Question 19

Zn2+(aq) + 2e- <=> Zn(s)
E(theta) = -0.76V
What will happen if the concentration of Zn2+ is decreased?

Answer 19

equilibrium shifts left, increasing the number of e- in the system
electrode potential becomes more negative

Question 20

What happens if standard electrode potentials are not measured at 1 moldm-3?

Answer 20

electrode potential will not be accurate: will increase/decrease depending on whether conc has increased/decreased

Question 21

What happens if the E(cell) is less than 0.06V?

Answer 21

reaction very unlikely to proceed

Question 22

Why may a reaction not be able to go ahead even if the electrode potential indicate that it could?

Answer 22

is thermodynamically feasible but reaction rate may be very slow i.e. catalyst needed

Question 23

How do changes in E(theta) due to concentration affect E(cell)?

Answer 23

may increase or decrease the difference between the 2 E(theta)s of the half cells
will alter E(cell) accordingly

Question 24

What is a primary cell?

Answer 24

non-rechargeable: reactions are not reversible: stops working when the chemical reacting at the electrodes run out
used in clocks
usually Zn and MnO2 reactants with KOH electrolyte

Question 25

What is a secondary cell?

Answer 25

rechargeable: reaction is reversible and reverse reaction occurs while recharging used in lithium ion batteries in phones

Question 26

How does a fuel cell work?

Answer 26

hydrogen rich fuel reacts w oxygen generating e- and a current fuel and oxygen flow into the cell, electrolyte allows transfer of ions + products flow out only water produced as product (in pure H2 cell)

Question 27

What are the advantages and disadvantages of H fuel cells?

Answer 27

AD: clean, quiet, 2/3x more efficient than burning fuel DIS: production involves toxic chemicals that need to be disposed of, cells are highly flammable

Question 28

What is the difference between using an alkaline vs acid electrolyte in a fuel cell?

Answer 28

in both cases: O2 in, fuel in, water out, E(cell) = 1.23V, there is a membrane separating half cells but have different redox systems

Question 29

What are the hydrogen oxygen fuel cell equations with an acidic electrolyte?

Answer 29

2H+ + 2e- <=> H2 : flip: becomes oxidation equation 1/2 O2 + 2H+ + 2e- <=> H2O : reduction overall: 1/2 O2 + H2 <=> H2O H+ travel to and water leaves from the oxygen electrode (cathode)

Question 30

What are the hydrogen oxygen fuel cell equations with an alkaline electrolyte?

Answer 30

2H2O + 2e- <=> H2 + 2OH- : flip: become oxidation equation 1/2 O2 + H2O + 2e- <=> 2OH- : reduction overall: 1/2 O2 + H2 <=> H2O OH- travel to and water leaves from the hydrogen electrode (anode)

Question 31

Practicals: iodine/thiosulphate redox titrations | What are the reduction, oxidation and overall equations for the titrations?

Answer 31

oxidation: 2(S2O3)2- (aq) => (S4O6)2- (aq) + 2e- reduction: I2(aq) + 2e- => 2I- (aq) overall: 2(S2O3)2- + I2 => (S4O6)2- + 2I-

Question 32

What is the I/THIO reaction used to analyse?

Answer 32

ClO- in bleach Cu2+ in copper (II) alloys Cu content in copper alloys (Cu converted to Cu2+ first)

Question 33

How do to I/THIO reactions work?

Answer 33

the chemical under analysis is reacted with excess KI to produce iodine (yellow-brown solution) excess = complete reaction solution titrated against Na2S2O3 and yellow-brown fades as I2 reduced back to I- when solution is pale straw colour, add starch indicator, turns blue-black when end point is reached, blue-black just disappears

Question 34

What happens if you add starch too early in an I/THIO titration?

Answer 34

solid black complex forms, titre is too low

Question 35

How would you prepare an insoluble copper sample for analysis?

Answer 35

react it with concentrated acid, then neutralise to form Cu2+ ions (then react w excess KI, titrate, etc.)

Question 36

What is the reducing agent in the I/THIO reactions?

Answer 36

(S2O3)2- = reducing agent when reacting with I2

Question 37

``` Practicals: manganate (VII) redox titrations What does (MnO4)- act as? ```

Answer 37

oxidising agent

Question 38

Why does (MnO4)- (and also dichromate) need to be acidified with excess H2SO4?

Answer 38

provides a source of H+ ions which allows the reaction to occur and provides an excess to prevent other products from forming

Question 39

What colour is manganate before and after reduction?

Answer 39

manganate: deep purple Mn2+ : very pale pink self indicating titration

Question 40

What is the end point of a manganate titration?

Answer 40

when the first permanent pale pink solution is reached

Question 41

What is KMnO4 used to analyse?

Answer 41

ethanedioic acid (oxidised to CO2)

Question 42

What is the colour change of (Cr2O7)2- when reduced?

Answer 42

(Cr2O7)2- : orange Cr3+ : green cannot self indicate, green is too strong

Question 43

Where do you read the meniscus from in manganate titrations?

Answer 43

the top, as the deep purple colour is too difficult to see through

Question 44

What is the process of a manganate titration?

Answer 44

dissolve compound and make up to 250cm^3 solution in volumetric flask add 25cm^3 to a conical flask w 10cm^3 H2SO4 ( an excess) fill burette with standard solution of KMnO4 + titrate

Question 45

redox: more positive = greater tendency to undergo...

Answer 45

reduction | non-metals tend to be more positive

Question 46

redox: more negative = greater tendency to undergo...

Answer 46

oxidation | metals tend to be more negative

Question 47

What are the limitations of electrode potential predictions?

Answer 47

if cell operates under non standard conditions, actual electrode potentials will be different to E theta values as the cell operates E theta values change activation energy could be too high or rate of reaction very slow

Question 48

Why is using a lower temperature for a reaction beneficial to the environment?

Answer 48

less fossils fuels/energy used reduction in CO2 emissions and global warming