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OCR Chemistry A > 5.2.3 - Redox & electrode potentials > Flashcards

5.2.3 - Redox & electrode potentials Flashcards

(57 cards)

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1
Q

Constructing half equations when the oxidation no. INCREASES

A

Delocalised e- goes on the products side

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2
Q

Constructing half equations when the oxidation no. DECREASES

A

Delocalised e- goes on the reactants side

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3
Q

Half eqns. in acidic conditions

A

Add H2Os to balance oxygens

H+ to balance H’s

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4
Q

Half eqns. in alkaline conditions

A

Add H2Os to balance oxygens

OH- to balance H’s

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5
Q

Combining half eqns to get the overall eqn

A

Balance delocalised e- then combine the two

Delocalised e- should cancel

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6
Q

Redox titrations

A

Involves transfer of e- from one species to another

Titrations of an oxidising agent against a reducing agent

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7
Q

Use of acidified KMnO4

A

Purple potassium manganate is in the burette
Sample analysed is in the flask with an excess of dilute sulfuric acid
As the MnO4- ions react, they form Mn 2+ ions which are colourless

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8
Q

Why can’t HCl be used w/ MnO4-

A

MnO4- would oxidise Cl- to Cl2 and then affect the vol of KMnO4 required in the titration

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9
Q

Why can’t conc. H2SO4/HNO3 be used w/ MnO4-

A

They are oxidising agents themselves so affect the vol of KMnO4 required

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10
Q

Why can’t ethanoic acid be used w/ MnO4-

A

It’s a weak acid and would not provide enough H+ ions

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11
Q

Reacting ratio of Fe 2+ : MnO4-

A

5:1

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12
Q

How is Fe (0) used in the redox titration

A

Oxidised w/ H2SO4 to Fe (2+) ready for analysis

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13
Q

How is Fe(+3) used in the redox titration

A

Reacted w/ Zn to reduce it to Fe (2+)

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14
Q

Reacting ratio of C2O4 2-: MnO4-

A

2.5 : 1

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15
Q

C2O4 2-

A

Ethanedioate ion

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16
Q

Self indicating

A

Titration that does not need an indicator

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17
Q

Autocatalysis

A

A reaction where one of the products acts as a catalyst

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18
Q

Reaction w/ MnO4- and C2O4 2-

A

As both ions are -ve they repel each other and the reaction is slow and needs warming at the start
However Mn2+ acts as a catalyst and speeds up reaction

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19
Q

Cu^2+/ I-

A

2 Cu2+ + 4I- —-> 2CuI + I2

CuI - white ppt

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20
Q

ClO-/I-

A

2 I- + ClO- + 2H+ —> Cl- + I2 + H2O

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21
Q

I2/S2O3 2-

A

2 S2O3 + I2 —> 2 I- + S4O6 2-

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22
Q

Reacting ratio Cu2+:I2

A

2:1

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23
Q

Reacting ratio of ClO- : I2

24
Q

Reacting ratio of S2O3 2- : I2

25
Electrochemistry
Control of transfer of e- to produce electrical energy
26
Standard electrode potential
The emf of a half cell compared to the standard hydrogen electrode, measured at 298 K w/ sol. conc. of 1 moldm3 and gas pressure of 101 kPa
27
Strong reducing agents have ...
more -ve electrode potentials and so get oxidised
28
Strong oxidising agents have ...
More +ve electrode potential and so get reduced
29
Calculating cell potential
More +ve electrode potential - more -ve electrode potential
30
Where does oxidation occur
Anode (+ve)
31
Where des reduction occur
Cathode (-ve)
32
How are half cells eqns written
Showing the reduction so e- are on the lhs | Species on the lhs are oxidising agents and species on the rhs are reducng agents
33
When is a cell reaction feasible
When the cell potential is +ve
34
Why may a feasible reaction not occur
HIgh Ea or conditions aren't standard/ diff conc
35
Explaining changes in a cell
..... affects which electrode The eqm shifts to the ... in order to .. ... accepts/ donates more e- The electrode potential is more +ve/-ve and the cell potential is more/less +ve
36
What is the salt bridge soaked in and why
KNO3 Highly soluble Easily enters solutions to balance charges Unlikely to form ppts
37
How do electrons travel in a cell
In the wire
38
How do ions travel in a cell
Through the salt bridge
39
Metal used in half cells
Platinum (s)
40
Features of primary cells
Non-rechargebale Can only be used once as the redox reactions arent reversible Chemicals get used up --> pd falls --> battery goes flat ---> discarded Used in low-current, long-storage devices e.g. wall clocks
41
Redox system in most primary cells
Zn/MnO2 (alkaline) Both involve H2O and OH- (cancel out) MnO2/ Mn2O3 - more +ve, OA, cathode Zn/ZnO - more -ve, RA, anode
42
Feautures of secondary cells
Rechargable - each reaction is reversed during recharging, regenerating chemicals Li + in modern devices e.g. phones, tablets Nickel/cadmium and nickel metal hydride in torches
43
Fuel cells
Uses energy from reactions w/ O2 to create a pd (combustion) Fuel and O2 flow into fuel cells and produts out. Electrolyte remains in cell Can operate continously and dont have to be recharged
44
Hydrogen fuel cells
Can be acid or alkali (generate same potential difference - same overall reaction) Form no CO2 during combustion, only H2O
45
Acid hydrogen fuel cells
H2 enters at anode H+ moves to cathode O2 enters at cathode H2O leaves at cathode
46
Alkali hydrogen fuel cells
H2 enters at anode H2O leaves at anode OH- moves to cathode O2 enters at cathode
47
Negative of hydrogen fuel cell
H2 is a gas and difficult to store
48
What's the purpose of a salt bridge
Separates the solution Complete the electrochemical circuit Balance charges by releasing K+ and NO3- into the different solutions
49
Identifying the +ve electrode using voltmeter readings
If +ve, electrode connected to +ve terminal (reduction eqn) | If -ve, electrode connected to -ve terminal (oxidation eqn)
50
Why does the emf of a cell change once connected to electrodes w/ a flowing current
Conc of ions change
51
Function of the platinum electrode
Allow transfer of electrons | Reaction surface
52
Why may a redox reaction not occur in the absence of light
High Ea | Light breaks bonds
53
Properties of platinum that makes it suitable as an electrode
Inert | Conductor
54
Reduction of H2O2
2H+ + H2O2 +2e- ----> 2H2O
55
pH of SHE
0
56
Difference between fuel cell and modern storage cell
Fuel reacts w oxygen to give electrical energy
57
Risks with Li based cells
Toxicity | Fire

OCR Chemistry A (16 decks)

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