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OCR A Level Chemistry > Chapter 23 - Redox and Electrode Potentials > Flashcards

Chapter 23 - Redox and Electrode Potentials Flashcards

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1
Q

What is an oxidising agent?

A

Takes electrons from species being oxidised. Contains the species that is being reduced

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2
Q

What is a reducing agent?

A

Adds electrons to the species being reduced. Contains the species that is being oxidised

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3
Q

What are some of the important oxidation numbers?

A

Uncombined element = 0
Combined hydrogen = +1
Combined oxygen = -2
Ion of element = ionic charge

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4
Q

What are the steps for constructing a redox equation from half equations?

A
  1. balance the electrons
  2. add and cancel electrons
  3. cancel any species that are on both sides of the equation
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5
Q

What are the steps for constructing a redox equation from oxidation numbers?

A
  1. assign oxidation numbers to identify atoms that change their oxidation number
  2. balance only the species that contain elements that have changed oxidation number
  3. balance any remaining atoms
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6
Q

Where do electrons appear in half equations for reduction?

A

Left hand side
e.g. Cl2 + 2e- → 2Cl-

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7
Q

Where do electrons appear in half equations for oxidation?

A

Right hand side
e.g. Cu → Cu2+ + 2e-

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8
Q

How can oxidation numbers be used to write a half equation?

A
  1. assign oxidation numbers and the change in oxidation number
  2. balance e- (e.g. decrease in oxidation number of 5 requires 5e- on left hand side)
  3. balance any remaining atoms and predict any further species.
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9
Q

What are the two common redox titrations?

A
  • Potassium manganate (VII) under acidic conditions
  • Sodium thiosulfate for determination of iodine
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10
Q

What are manganate titrations used for the analysis of?

A

Reducing agents e.g. iron ions or ethanedioic acid

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11
Q

How does manganate change in redox titrations?

A

Reduced from MnO4 - to Mn2+

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12
Q

What are iodine-thiosulfate titrations used for the analysis of?

A

Oxidising agents e.g. Cu ions or chlorate ions

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13
Q

What is oxidised and what is reduced in iodine-thiosulfate reactions?

A

Thiosulfate ions are oxidised
Iodine is reduced

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14
Q

What are the 2 half equations for iodine-thiosulfate reactions?

A

Oxidation = 2S2O3 2- → S4O6 2- + 2e-
Reduction = I2 +2e- → 2I-

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15
Q

What is the overall equation for iodine-thiosulfate reactions?

A

2S2O3 2- + I2 → 2I- + S4O6 2-

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16
Q

What is the method for iodine-thiosulfate titrations?

A
  1. excess of acidified iodide ions and known volume of oxidising agent react in a conical flask
  2. products of reaction 1 left in a conical flask, titrated with known conc of sodium thiosulfate
  3. add starch when the solution has turned straw yellow, the solution will go blue black and the endpoint is when the conical flask goes colourless
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17
Q

What is an electrochemical cell?

A

A voltaic cell which converts chemical energy into electrical energy via the transfer of electrons

18
Q

What is a half cell?

A

Contains the chemical species present in a redox half equation

19
Q

What is a metal/metal ion half cell?

A

A simple half cell made up of a metal rod dipped into a solution of its aqueous metal ion. This sets up an equilibrium
M 2+ + 2e ⇌ M (s)
One metal half cell on its own will have no net electron transfer.

20
Q

What is an ion/ion half cell?

A

A half cell containing ions of the same element in different oxidation states. This sets up a redox equilibrium:
Fe3+ + e- ⇌ Fe2+
Uses a platinum electrode to transport electrons in and out of the half cell

21
Q

When two metal/metal ion half cells are connected, how do you know where oxidation and reduction occur?

A

Oxidation happens at the more reactive electrode (e- are more easily lost)
Reduction happens at the less reactive electrode (e- more easily gained)
This creates an electrode potential difference

22
Q

What is the standard electrode potential?

A

The e.m.f. of a half cell connected to a standard hydrogen half cell under standard conditions (state standard conditions!!! 1mol dm-3 1 atm/100kPa and 298K)

23
Q

How is a standard hydrogen electrode set up?

A

Acid solution containing 1.0 mol dm-3 H+ with platinum electrode inserted. Glass tube with holes in it to allow bubbles of H2 gas to escape. H2 at 298K and 100kPa

24
Q

How do you connect 2 half cells?

A
  • electrodes connected by a wire to allow a controlled flow of electrons
  • two solutions connected with a salt bridge to allow ions to flow
25
Q

What is a salt bridge?

A

Typically a concentrated solution of an electrolyte which doesn’t react with either solution e.g. filter paper soaked in aq KNO3

26
Q

What does the E o value tell you?

A

More negative = greater tendency to lose electrons and undergo oxidation
More positive = greater tendency to gain electrons and undergo reduction

27
Q

How do you calculate standard electrode potential?

A

E cell = E (positive electrode) - E (negative electrode)

28
Q

How can you predict if a redox reaction is feasible?

A

Using the electrochemical series (standard electrode potentials)

29
Q

What does it tell you if a redox system has a more negative standard E cell value?

A

More negative = greater tendency to be oxidised and LOSE electrons
Therefore better REDUCING AGENTS

30
Q

What does it tell you if a redox system has a more positive standard E cell value?

A

More positive = greater tendency to be reduced and GAIN electrons
Therefore better OXIDISING AGENTS

31
Q

How do you connect two redox systems from the electrochemical series?

A

Connect the two with an anti clockwise circuit (more negative redox system on the top)
- keep more positive redox system the SAME
- CHANGE the more negative redox system

32
Q

How does changing concentration affect standard electrode potentials?

A

If concentration of solution is not 1 mol dm-3, then value of the electrode potential will be different from standard value
Shift to the side with less electrons makes electrode potential less negative
Shift to side with more electrons makes electrode potential more negative

33
Q

What are the 3 types of cell?

A

Primary, secondary and fuel

34
Q

What are primary cells?

A

Cells which are non renewable and designed to be used once only. When in use, electrical energy is produced by oxidation and reduction at the electrodes.
Most modern primary cells = Zn/MnO2 and KOH alkaline electrolyte

35
Q

What are some examples where primary cells are used?

A

Low current, long storage devices like wall clocks or smoke detectors

36
Q

What are secondary cells?

A

Rechargeable cells, where the cell reaction produced electrical energy can be reversed during recharging.
Common examples = nickel-metal hydride or lead-acid

37
Q

What are some examples where secondary cells are used?

A

Car batteries, radios and torches, laptop and phone batteries

38
Q

What are fuel cells?

A

A fuel cell uses the energy from the reaction of a fuel with oxygen to create a voltage.
Fuel and oxygen flow into fuel cell, products flow out. Can operate continuously provided fuel and oxygen are supplied, do not have to be recharged
Common examples = hydrogen fuel cells or methanol

39
Q

What are the advantages and disadvantages of fuel cells over fossil fuels?

A

Advantages: only H2O produced/non polluting, greater efficiency
Disadvantages: hydrogen gas is difficult to store, and difficult to manufacture initially

40
Q

What is the flow of electrons in 2 half cells?

A

Flows from more negative to half cell with more positive standard electrode potential

OCR A Level Chemistry (29 decks)

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