23) Redox and electrode potentials Flashcards

1
Q

What two things will there always be in a redox reaction?

A

oxidising agent

reducing agent

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2
Q

Define oxidising agent

A

a reagent that oxidises (takes e- from) another species

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3
Q

Define reducing agent

A

a reagent that reduces (adds e- to) another species

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4
Q

What should you do when writing an overall equation for redox reactions?

A

balance the changes in oxidation number and then balance any remaining atoms afterwards

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5
Q

If one species in an overall redox equation has an oxidation number change of +6, and a second species has an oxidation number change of -1, what should be done?

A

multiply second species by 6 to give -6

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6
Q

Describe manganate (VII) titrations

A

a standard solution of KMnO4 is added to burette with excess of dilute sulfuric acid
as it reacts it is decolourised from deep purple colour until end point = permanent pink colour - indicating an excess of MnO4 - ions (self-indicating)

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7
Q

What is the purpose of excess of dilute sulfuric acid in a manganate (VII) titration?

A

provides H+ for the reduction of MnO4 - ions

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8
Q

Why are burette readings taken from the top of the meniscus in a manganate (VII) titration?

A

manganate (VII) solution is a deep purple colour

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9
Q

What can manganate (VII) titrations be used for?

A
  • analysis of reducing agents e.g. Fe 2+ (analysing the purity of an iron (II) compound) or ethanedioic acid (COOH)2
  • reduce MnO4 - to Mn 2+
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10
Q

What can iodine/thiosulfate titrations be used to determine?

A
  • chlorate ions, ClO- content in household bleach (HCl added to acidify solution)
  • Cu 2+ content in copper (II) compounds
  • Cu content in copper alloys (dissolved in conc. nitric acid + neutralisation to produce Cu 2+)
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11
Q

Describe iodine/thiosulfate titrations

A
  • a solution of oxidising agent to be analysed is added to a conical flask with an excess of potassium iodide - iodine produced resulting in a yellow-brown colour
  • titrated with sodium thiosulfate solution (S2O3 2- oxidised and I2 reduced)
  • when end point is approached, sol. has faded to a pale straw colour, a small amount of starch indicator is added -> blue-black solution which turns colourless at end point
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12
Q

Define voltaic cell

A

a type of electrochemical cell which converts chemical energy into electrical energy

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13
Q

What does a half cell contain?

A

the chemical species present in a redox half-equation

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14
Q

How is a voltaic cell made?

A

by connecting together 2 different half-cells allowing e- flow

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15
Q

Describe metal/metal ion half-cells

A

a metal rod dipped into a solution of its aqueous metal ion
at the phase boundary where the metal is in contact with its ions, an equilibrium is set up e.g. Zn 2+(aq) + 2e- <=> Zn(s)

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16
Q

Describe ion/ion half-cells

A

contains ions of the same element in different oxidation states
there is no metal to transport e- into/out of the half-cell so an inert metal electrode made out of platinum is used

17
Q

What happens in an operating cell with two metal/metal ion half-cells connected?

A

the electrode with more reactive metal loses e- and is oxidised
the electrode with the less reactive metal gains e- and is reduced

18
Q

Define standard electrode potential E⦵

A

the emf of a half-cell compared with a standard hydrogen half-cell under standard conditions

19
Q

What is the standard electrode potential of a standard hydrogen half-cell under standard conditions?

A

0V

20
Q

How is standard electrode potential measured?

A

the half-cell is connected to a standard hydrogen electrode by a wire to allow controlled flow of e-
recorded by voltmeter

21
Q

Describe a standard hydrogen half-cell

A

platinum electrode
H2(g) pumped in
acid sol. 1moldm-3 H+(aq)

22
Q

What does a salt bridge typically contain?

A

a concentrated solution of electrolyte that does not react with either solution e.g. a strip of filter paper soaked in KNO3(aq)

23
Q

How is the equilibrium for standard electrode potentials shown?

A

with the forward reaction being the reduction

24
Q

How are electrode potentials sorted in a data reference table?

A

in order of E⦵ - most negative at top

25
Q

The more negative E⦵ value, the greater the tendency to _ e- and undergo _?

A

lose

oxidation

26
Q

The more positive E⦵ value, the greater the tendency to _ e- and undergo _?

A

gain

reduction

27
Q

E⦵cell = ?

A

E⦵(positive electrode) - E⦵(negative electrode)

28
Q

What should you remember when predicting redox reactions from a data reference table?

A

top-right, bottom-left rule
the strongest reducing agent is at the top right
the strongest oxidising agent is at the bottom left

29
Q

What should you remember when writing an overall equation from half-equations in a data reference table?

A

the oxidation half-equation is obtained by reversing the equilibrium

30
Q

What are the limitations of predicting redox reactions from a data reference table? (4)

A
  • gives no indication of rate of a reaction
  • if concentration is not equal to 1moldm-3 electrode potential will be different to standard value
  • may not be under standard conditions
  • many reactions are not aqueous (aqueous equilibria)
31
Q

Describe primary cells

A
  • non-rechargeable
  • electrical energy is produced by reduction and oxidation at electrodes (irreversible reactions)
  • eventually chemicals will be used up, voltage decreases and battery goes flat
  • alkaline (Zn/MnO4 and potassium hydroxide alkaline electrolyte, Zn oxidised -> ZnO and Mn in MnO2 reduced -> Mn2O3)
32
Q

Describe secondary cells

A
  • rechargeable - cell reaction can be reversed during recharging so chemicals are regenerated
  • lithium-ion and lithium ion polymer cells (when in use e- flow from negative -> positive electrode and Li+ from anode to cathode to keep electrical charge balance)
33
Q

Describe fuel cells

A
  • uses the energy from the reaction of a fuel with oxygen to create voltage
  • fuel and oxygen flow into cell and products flow out, electrolyte remains
  • operate continually provided fuel and O2 supplied
  • do not have to be recharged
34
Q

Give the overall equation for hydrogen fuel cells and the E⦵cell

A

H2(g) + 1/2 O2(g) -> H2O(l)

E⦵cell = 1.23V

35
Q

Suggest two advantages of using methanoic acid as the fuel in a fuel cell rather than hydrogen

A

liquid - easier to store and transport
greater cell potential / voltage
hydrogen is flammable