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A Level Chemistry > Module 5: Chapter 23 - Redox and Electrode Potentials > Flashcards

Module 5: Chapter 23 - Redox and Electrode Potentials Flashcards

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1
Q

What is the charge of CO?

A

0

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2
Q

What is the oxidation state of Nickel in Ni(CO)₄?

A

0

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3
Q

What is the thiosulfate ion?

A

S₂O₃²⁻

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4
Q

Why is the thiosulfate ion, S₂O₃²⁻, commonly used in redox titrations?

A

It is a reducing agent

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5
Q

Explain the oxidation numbers in S₄O₆²⁻

A

The oxidation numbers of all oxygen atoms is -2. Therefore, it would appear that the oxidation number on each sulfur atom is +2.5. However, it cannot actually be +2.5 as you cannoy have have an electron, therefore the true case would be 2 of the S atoms having an oxidation numberof +2 and 2 of the S atoms having an oxidation number of +3. However, for the sake of calculations and writing half equations it can be treated as +2.5

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6
Q

What is a strong oxidising agent?

A

Potassium Manganate (VII), KMnO₄

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7
Q

What is the half equation for Manganate (VII) being reduced to manganese (II) ions?

A

MnO₄⁻ + 5e⁻ + 8H⁺ ⇌ Mn²⁺ + 4H₂O

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8
Q

What are the conditions to use potassium managate (VII) in redox titrations?

A

It must be in acidic conditions to provide the H⁺ ions. Specifically dilute sulfuric acid must be used in excess

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9
Q

Why must dilute sulfuric acid be used with potassium managate (VII) in redox titrations rather than hydrochloric acid, concentrated sulfuric/nitric acid, or ethanoic acid?

MnO₄⁻ + 5e⁻ + 8H⁺ ⇌ Mn²⁺ + 4H₂O

A
  • Hydrochloric acid - Manganate will oxidise the chloride to form chlorine aswell as the target. This will cause the titre to increase
  • Concentrated sulfuric/nitric acid - They are oxidising agents so will decrease Mn²⁺ produced, titre will decrease
  • Ethanoic acid - It is a weak acid so insufficient H⁺ ions will be put into the solution
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10
Q

What is a self-indicating titration?

A

A titration where no indicator is used as the solution changes colour during the reaction

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11
Q

What colour is potassium manganante?

A

Purple

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12
Q

What colour are manganese (II) ions?

A

Colourless
(Technically very pale pink)

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13
Q

How can you determine the end point in a redox titration using potassium manganate (VII) as an oxidising agent?

A

The purple potassium manganate is in the burette and is titred into the colourless reducing agent. As the potassium manganate is titred in, it will form colourless manganese (II) ions, however once the titration is complete and all of the reducing agent has been oxidised, this will no longer happen. Therefore the end point is when the solution turns light pink

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14
Q

How can the element iron be prepared to be in the +2 oxidation state?

A

If is it the element Fe (0), it is reacted with sulfuric acid to oxidise it to form Fe²⁺

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15
Q

How can Fe³⁺ be prepared to be in the +2 oxidation state?

A

If it is Fe³⁺, it is reacted with Zn to reduce it to Fe²⁺

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16
Q

What is the overall equation for the redox titration between Fe²⁺ and acidified potassium manganate (VII)

A

MnO₄⁻ + 5Fe²⁺ + 8H⁺ ⇌ Mn²⁺ + 5Fe³⁺ + 4H₂O

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17
Q
A

94%

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18
Q

What is a half-cell?

A

An electrode

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19
Q

Explain how a metal half-cell/electrode works

A
  • When a piece of metal is dipped into a solution of its own metal ions, an equilibrium is set up. There is a tendency for the metal to form positive ions and go into solution, however there is also a tendency for the metal ions in solutions to gain electrons and form metal. Mⁿ⁺(aq) + ne⁻ ⇌ M(s)
  • If this equilibrium lies to the left, then the metal aquires a negative charge due to a build up of electrons on the metal
  • If this equilibrium lies to the right, then a positive charge builds up on the metal as electrons have been used up to form the solid metal from the ions.
  • The position of equilibrium (and so the charge of the electrode) depends on the reactivity of the metal
  • A more reactive metal tends to form Mⁿ⁺ ions and is therefore more likely to be negative. A less reactive metal is more likely to have a positive charge
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20
Q

What determines if an electrode is positive or negative?

IMPORTANT

A

Its reactivity compared to the other electrode:
* If the electrode is more reactive, then it is a better reducing agent and is therefore oxidised. As a result it is better at releasing electrons (which remain in the electrode) and therefore forms the ions and gives the electrode a negative charge. Therefore it is the negative electrode.
* If the electrode is less reactive, then it is a worse reducing agent (better oxidising agent) and is therefore reduced. As a result it is worse at releasing electrons and therefore gains the electrons to form the elements (giving the electrode a positive charge). Therefore it is the positive electrode

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21
Q

Is the more reactive metal electrode positive or negative?

A

Negative electrode

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22
Q

Is the less reactive metal electrode positive or negative?

A

Positive electrode

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23
Q

What are the 3 types of electrodes?

A
  • Metal electrodes
  • Gas electrodes
  • Redox electrodes
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24
Q

Explain how a Gas electrode works

A

A gas electrode is for a gas and a solution of its ions. An inert metal (usually platinum) is the actual electrode to allow the flow of electrons

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25
Q

Explain how a redox electrode works

A

A redox electrode is for 2 different ions of the same element where the 2 types of ions are present in solution with an inert metal electrode (usually platinum) to allow the flow of electrons

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26
Q

How can you measure the potential of an electrode?

A

The actual potential of an electrode cannot be directly measured, therfore it is connected to another half-cell of known potential and the potential difference between the 2 half-cells is measured. Therefore a standard hydrogen electrode (SHE) is assigned the potential of 0V and all the other halfcells can be measured against it.

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27
Q

What does combining 2 half-cells produce?

A

An electrochemical cell

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28
Q

What is the emf of a Standard hydrogen electrode?

A

0V

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29
Q

How do you set up an electrochemical cell?

Include diagram

A

2 Half cells are joined together to give a complete circuit
* The 2 metals are joined with a wire (electrons flow through the wire)
* The 2 solutions are joined with a salt bridge (ions flow through the salt bridge)
* A voltmeter is often included in the circuit to allow the potential difference (emf) to be measured

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30
Q

What does a salt bridge consist of?

A

Either:
* A piece of filter paper soaked with a solution of unreactive ions
* A tube containing unreactive ions in agar gel

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31
Q

What compound is commonly used in a salt bridge?

A

Potassium Nitrate as potassium ions and nitrate ions are quite unreactive

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32
Q

What are the standard conditions when measuring the potential of a half-cell?

A
  • Cell concentration - 1.0 mol dm⁻³ of the ions invovled in the half equation
  • Cell temperature - 298K
  • Cell pressure - 100kPa

It must also be measured under 0 current conditions in order to measure the full emf (this can be achieved with a high resistance voltmeter)

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33
Q

How would changing the concentration of the ionic solution change the emf of the half-cell?

A

Mⁿ⁺(aq) + ne⁻ ⇌ M(s), by changing the concentration of Mⁿ⁺, the position of equilibrium would move making the potential more or less negative

34
Q

What is the standard electrode potential, Eθ?

A

The emf of a half cell connected to a standard hydrogen half cell under standard conditions of 298K, solution concentrations of 1 mol dm⁻³, and a pressure of 100kPa

35
Q

Draw the diagram for a standrd hydrogen electrode

A
36
Q

What must the solution be in a standard hydrogen electrode?

A

It must be a monoprotic acid

37
Q

What is the primary standard?

A

The standard to which all other electrode potentials are compared (a SHE)

38
Q

What is the cell notation of a standard hydrogen electrode?

A

Pt(s) | H₂(g) | H⁺(aq)

39
Q

Draw an ion, ion half cell

A
40
Q

What is the electrochemical series?

A

A list of standard electrode potentials in order of decreasing/increasing potential. It matches the reactivity series with the elements which are the strongest reducing agents having the most negative values, and the elements which are the worst reducing agents having the most positive values.

41
Q

How can you calculate the cell emf?

A

Cell emf = Eθ(positive terminal) - Eθ(negative terminal)

42
Q

How can you construct the redox equation for 2 half cells?

A
  1. Change the half equation for the more negative Eθ value in order to show oxidation (it releasing electrons)
  2. Balance the electrons
  3. Combine 2 half equations and cancel electrons
43
Q

Construct the equation for the spontaneous reaction and calculate the electrochemical cell emf

A

Fe(s) + 2Ag⁺(aq) ⇌ 2 Ag(s) + Fe²⁺(aq)

emf = +1.24

44
Q

Construct the equation for the spontaneous reaction and calculate the electrochemical cell emf

A

S₂O₈²⁻(aq) + 2Ag(s) ⇌ 2Ag⁺(aq) + 2SO₄²⁻

emf = 1.21V

45
Q

How would changing the pressure of the gas change the emf of the gas half-cell?

A

Changing the pressure can cause the position of equilibrium to move, therefore the number of electrons being given out/in can change and will effect the emf of the half cell

46
Q

How would changing the temperature of the system change the emf of the half cell?

A

Changing the temperature can cause the position of equilibrium to move, therefore the number of electrons being given out/in can change and will effect the emf of the half cell

47
Q

How would using a piece of platinum with a greater surface area change the emf of a gas half cell?

A

It would have no effect, it would allow the rate of electron exchange to increase, but it would have no effect on the potential

48
Q

How would using sulfuric acid at half the concentration of hydrochloric acid effect the emf of the half cell?

A

It wouldnt effect the emf of the half cell as the concentration of hydrogen ions would remain the same

49
Q

What are non-rechargable cells (Primary cells)?

A

In non-rechargeable cells, the chemicals are used up over time and the emf drops. Once one or more of the chemicals has been completely used up, the emf drops to 0 volts. They cannot be recharged and they have to be disposed of after their single use

50
Q

What are 2 types of non-rechargable cells?

A
  • Zinc-carbon cell
  • Alkaline Cell
51
Q

What is a zinc-carbon non-rechargeable cell?

A

It is the standard, cheap non-rechargeable cell. It is cheap but has a fairly short life

52
Q

What is an alkaline non-rechargeable cell?

A

It is a higher cost cell but has a longer life

53
Q

What are primary cells?

A

Non-rechargeable cells

54
Q

What are rechargeable cells (Secondary cells)?

A

In rechargeable cells the reactions are reversible, they are reversed by applying an external current which regenerates the chemicals

55
Q

What are the 3 types of rechargeable cells?

A
  • Lithium ion cells
  • Lead-acid cells
  • Nickel-cadmium cells
56
Q

What is a secondary cell?

A

A rechargeable cell

57
Q

When are lithium ion cells used?

A

They are used in phones, tablets, cameras, laptops, etc

58
Q

When are lead-acid cells used?

A

In car batteries

59
Q

When are nickel-cadmium batteries used?

A

They are cylindrical batteries used in radios, torches. Same uses as non-rechargable cells but they can be recharged

60
Q

What are 2 common redox titrations?

A
  • Potassium Manganate (VII) under acidic conditions
  • Sodium thiosulfate for determination of iodine
61
Q

What is the thiosulfate ion?

A

S₂O₃²⁻

62
Q

What is a voltaic cell?

A

A cell which converts chemical energy into electrical energy

63
Q

What are the limitations of making predictions using Eθ values?

A
  • Reaction rate - Although the electrode potential may indicate the thermodynamic feasiblity of a reaction, they give no indication of the rate of reaction
  • Conditions- The actual conditions may be different to that of the standard conditions, this will affect the vaue of the electrode potential
64
Q

How can you predict whether a potential redox reaction is feasible?

A

By using standard cell potentials, the reducing agent (gets oxidised) should have a more negative Eθ. The oxidising agent (gets reduced) should have a more positive Eθ

65
Q

What are 2 types of redox titrations that can be used for analysis?

IMPORTANT

A
  • Manganate (VII) redox titrations
  • Iodine/thiosulfate redox titrations
66
Q

Explain how an iodine/thiosulfate redox titrationt takes place and can be used to analyse oxidising agents?

A
  1. First the oxidising agent being investigated must react with Iodide and hydrogen ions to form iodine
  2. Then perform the titration by adding sodium thiosulfate
  3. You can then calculate the initial concentration of the oxidising agent being investigated
67
Q

What is the iodine/thiosulfate redox reaction?

A

The thiosulfate ions are oxidised and the iodine is reduced:
2S₂O₃²⁻(aq) + I₂(aq) -> 2I⁻(aq) + S₄O₆²⁻(aq)

68
Q

What is the indicator in iodine/thiosulfate redox reactions?

A

Once the end point is being approached and the iodine coour has faded enough to become a pale straw colour, add a small amount of starch indicator. The blue-black colour should fade once all the iodine has been reduced to iodide

69
Q

Explain how an iodine/thiosulfate redox titrationt takes place and can be used to analyse copper?

A
  1. First the copper alloy being investigated must react with concentrated nitric acid to form Cu²⁺ ions
  2. Then the ions must react with Iodide ions to form iodine (and a white precipitate of copper iodide)
  3. Then perform the titration by adding sodium thiosulfate
  4. You can then calculate the initial mass of copper being investigated
70
Q

How are fuel cells different from traditional cells?

A

Fuel cells require a continuous supply of chemicals into the cell, therefore they can operate continuously and do not run out of chemicals or need recharging. But they do require a constant supply of chemicals

71
Q

What is the most common type of fuel cell?

A

Hydrogen-oxygen fuel cell

72
Q

How does a hydrogen-oxygen fuel cell work?

A

The hydrogen and oxygen flow into the fuel cell and the products flow out. The energy of this reaction is used to create a voltage

73
Q

What are the emissions of a hydrogen fuel cell?

A

They produce no carbon emissions, only water

74
Q

What are the 2 types of hydrogen-oxygen fuel cells and how do they differ?

A
  • Alkaline
  • Acidic

They are exactly the same, however in an alkanline cell there is a movement of hydroxide ions in the electrolyte, and in the acidic cell there is a movement of hydrogen ions in the electrolyte

75
Q

What is the overall equation for hydrogen fuel cells?

A

H₂(g) + ½O₂(g) -> H₂O(l)

It is the same for both alkali and acidic cells

76
Q

What is the emf of a hydrogen fuel cell?

A

1.23V

It is the same for both alkali and acidic cells

77
Q

Describe the structure of an alkaline hydrogen fuel cell and the half equations at each electrode

A
78
Q

Describe the structure of an acidic hydrogen fuel cell and the half equations at each electrode

A
79
Q

Explain how you can make predictions about the feasibility of redox reactions from electrode potentials:

IMPORTANT

A

A redox reaction is only feasible if:
* The standard electrode potential of the reducing agent (species being oxidised) is the most negative in the system
* The standard electrode potential of the oxidising agent (species being reduced) is the most positive in the system

80
Q

What 2 pieces of data can be used to make predictions about the thermodynamic feasibility of a reaction?

A
  • Gibbs Free energy
  • Standard electrode potentials

A Level Chemistry (31 decks)

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