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A Level Chemistry > 23. Redox and Electrode Potentials > Flashcards

23. Redox and Electrode Potentials Flashcards

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1
Q

How do you balance half equations in basic solutions?

A

balance thing being oxidised/reduced

deduce e- needed from oxidation state change

balance on other side w OH-

balance on other side w H2O

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2
Q

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

A

chemical reactions that transfer e- from one species to another

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3
Q

When is a platinum electrode used?

A

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-

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4
Q

What does the standard electrode potential represent?

A

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

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5
Q

Define standard electrode potential.

A

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

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6
Q

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

A

Page 3 of electrode potentials leaflet

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7
Q

What is a salt bridge and what is its purpose?

A

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

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8
Q

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

A

high resistance voltmeter: measures potential for electron transfer

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9
Q

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

A

reduction half equations

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10
Q

What does decreasing E(theta) mean?

A

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

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11
Q

What does increasing E(theta) mean?

A

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

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12
Q

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

A

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

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13
Q

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

A

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

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14
Q

At which electrode does oxidation take place?

A

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

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15
Q

At which electrode does reduction take place?

A

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

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16
Q

At which electrode are electrons produced?

A

the electrode where oxidation is taking place: the negative electrode

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17
Q

Which direction do electrons flow?

A

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

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18
Q

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

A

equilibrium shifts right, removing e- from the system

electrode potential becomes less negative

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19
Q

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

A

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

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20
Q

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

A

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

21
Q

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

A

reaction very unlikely to proceed

22
Q

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

A

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

23
Q

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

A

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

24
Q

What is a primary cell?

A

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

25
Q

What is a secondary cell?

A

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

26
Q

How does a fuel cell work?

A

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)

27
Q

What are the advantages and disadvantages of H fuel cells?

A

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

28
Q

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

A

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

29
Q

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

A

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)

30
Q

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

A

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)

31
Q

Practicals: iodine/thiosulphate redox titrations

What are the reduction, oxidation and overall equations for the titrations?

A

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

32
Q

What is the I/THIO reaction used to analyse?

A

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

33
Q

How do to I/THIO reactions work?

A

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

34
Q

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

A

solid black complex forms, titre is too low

35
Q

How would you prepare an insoluble copper sample for analysis?

A

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

36
Q

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

A

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

37
Q 
Practicals: manganate (VII) redox titrations
What does (MnO4)- act as?
A

oxidising agent

38
Q

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

A

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

39
Q

What colour is manganate before and after reduction?

A

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

40
Q

What is the end point of a manganate titration?

A

when the first permanent pale pink solution is reached

41
Q

What is KMnO4 used to analyse?

A

ethanedioic acid (oxidised to CO2)

42
Q

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

A

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

43
Q

Where do you read the meniscus from in manganate titrations?

A

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

44
Q

What is the process of a manganate titration?

A

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

45
Q

redox: more positive = greater tendency to undergo…

A

reduction

non-metals tend to be more positive

46
Q

redox: more negative = greater tendency to undergo…

A

oxidation

metals tend to be more negative

47
Q

What are the limitations of electrode potential predictions?

A

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

48
Q

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

A

less fossils fuels/energy used

reduction in CO2 emissions and global warming

A Level Chemistry (30 decks)

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