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A Level Chemistry Year 2 > Redox and Electrode Potentials > Flashcards

Redox and Electrode Potentials Flashcards

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1
Q

What is a loss of electrons called?

A

Oxidation

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2
Q

What is a gain of electrons called?

A

Reduction

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3
Q

What is a redox reaction?

A

A reaction where oxidation and reduction happen simultaneously

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4
Q

What is an oxidising agent?

A

A substance that accepts electrons and gets reduced

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5
Q

What is a reducing agent?

A

A substance that donates electrons and gets oxidised

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6
Q

What is a redox reaction made up of, and how can you represent this?

A

Made up of an oxidation half-reaction and a reduction half-reaction, which you can write both as ionic half-equations

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7
Q

Write an ionic half-equation for the oxidation of iron

A

Fe —> Fe 3+ + 3e-

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8
Q

Write an ionic equation for the reduction of oxygen

A

O2 + 4e- —> 2O 2-

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9
Q

How do you work out the overall redox equation?

A

Write out the 2 half-equations, then make sure they have the same number of electrons, and combine and cancel out

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10
Q

How can you balance out half-equations if the oxidising/reduction agent contains oxygen or hydrogen?

A

You can add water, H+ ions (if acidic conditions) or OH- ions (if alkaline conditions)

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11
Q

What is the oxidation number of an element?

A

It tells you the total number of electrons it has donated or accepted

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12
Q

What does it mean if a chemical name has roman numerals after it?

A

It tells you the oxidation number of the element e.g. dichromate (VI) has oxidation number +6

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13
Q

What happens to the oxidation number for each electron lost?

A

It increases by 1

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14
Q

What happens to the oxidation number for each electron gained?

A

It decreases by 1

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15
Q

If the oxidation number is reduced, it’s…?

A

Reduction

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16
Q

What are concordant results?

A

Readings that are within a small range of each other

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17
Q

Why are transition metal ions good oxidising or reducing agents?

A

Because of their ability to gain or lose electrons easily

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18
Q

Why is it easy to spot when a transition element changes oxidation number?

A

Because a colour change occurs

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19
Q

What is the purpose of doing titrations using transitions elements irons?

A

You can find out how much oxidising agent is needed to exactly react with a quantity of reducing agent

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20
Q

What are redox titrations?

A

A titration that can be done to find out how much oxidising agent is needed to exactly react with a quantity of reducing agent, or vice versa

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21
Q

When doing redox titrations, why do you add an excess amount of dilute sulfuric acid to the conical flask containing a known amount of reducing agent?

A

To make sure there are enough H+ ions to allow the oxidising agent to be reduced

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22
Q

What do you do at the very start of a redox titration?

A

Measure out a quantity of reducing agent, using a pipette, then place in a conical flask

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23
Q

For a redox titration, what do you do once you’ve added the excess amount of dilute sulfuric acid to the conical flask?

A

You add some oxidising agent to a burette and take an initial reading of the volume, then gradually add the oxidising agent to the reducing agent

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24
Q

For a redox titration, what happens when you’re adding the oxidising agent to the reducing agent?

A

The oxidising agent you’re adding will react with the reducing agent, this reaction will continue until all of the reducing agent is used up

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25
Q

How do you know when the end point is reached for a redox titration?

A

The next drop of oxidising agent you add will result in a colour change to the colour of the oxidising agent

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26
Q

What do you do in a redox titration when a colour change in noticed?

A

Record the value for the volume of the oxidising agent added

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27
Q

For a redox reaction, what do you do differently between the rough redox titration and the accurate redox titration?

A

During the rough titration, you add the oxidising agent to the flask until the solution becomes a very faint colour of the oxidising agent. For the accurate titration, you add the oxidising agent to within 2cm^3 of the rough titration, then add drop by drop until the colour changes

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28
Q

How many times should you repeat doing an accurate redox titration?

A

Until you get 3 readings which are within 0.10cm^3 of each other

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29
Q

For a redox titration, how many decimal places should you record the volume added (titre)?

A

2 decimal places ending with either a 0 or a 5

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30
Q

For any titration, how do you measure the titre using a burette?

A

Measure from the bottom of the meniscus, and at eye-level

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31
Q

When doing the redox titration with KMnO4, how do you know when the endpoint is reached?

A

MnO4- ions in aqueous KMnO4 are purple, so the endpoint is when the solution has a colour change from colourless to purple

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32
Q

Give an example of a way you can see whether a colour change occurs for a redox titration?

A

You could hold a piece of paper behind the flask or put the flask on a white tile to see easier when the colour starts to change

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33
Q

What are the 4 steps to calculate the concentration of one of the reagents in a redox titration?

A

1- Write out a balanced redox equation for the reaction in the flask
2- For the reagent you know both the concentration ad volume for, calculate the number of moles
3- Use the molar ratios in the balanced equation to work out the number of moles of the reagent you want to find the concentration of
4- Calculate the unknown concentration using equation conc = (moles x 1000) /volume (cm^3)

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34
Q

What are iodine-sodium thiosulfate titrations used to find?

A

The amount of iodine in solution

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35
Q

What are the 4 steps for the iodine-sodium thiosulfate titration?

A

1- Oxidise the iodine ions to iodine
2- Titrate the iodine solution with sodium thiosulfate
3- Calculate the number of iodine moles present
4-Calculate the concentration of the oxidising agent

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36
Q

For the iodine-sodium thiosulfate titration, what chemical do you use as the oxidising agent?

A

Potassium iodate (V)

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37
Q

For the iodine-sodium thiosulfate titration, how do you oxidise the iodine ions to iodine?

A

Measure a certain volume of oxidising agent (Potassium iodate (V)). Add this to an excess amount of acidified potassium iodide solution. The iodate ions oxidise some of the iodide ions to iodine

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38
Q

What is the ionic equation for the oxidising of iodide ions to iodine?

A

IO3 - + 5I- + 6H+ —> 3I2 + S4O6 2-

39
Q

What are the 5 steps for the titration method for the iodine-sodium thiosulfate titration?

A
  • Take a flask containing the solution made when oxidising the iodide ions
  • From a burette, add sodium thiosulfate solution to the flask drop by drop
  • When the iodide colour fades away, add 2cm^3 of starch to detect the iodine. The solution will go dark blue showing there is some iodine still left
  • Add sodium thiosulfate one drop at a time until the blue colour disappears
  • Record the volume of sodium thiosulfate added to the solution
40
Q

What is the colour of the solution in the conical flask at the start of an iodine-sodium thiosulfate titration?

A

Red-brown

41
Q

What is the equation for the iodine-sodium thiosulfate titration?

A

I2 + 2S203 2- —> 2I- + S4O6 2-

42
Q

What is an electrochemical cell?

A

An electrical circuit made from 2 metal electrodes (connected by a wire) which are dipped in salt solution (connected by a salt bridge)

43
Q

What is an electrochemical series?

A

A list of electrode potentials written from most negative to most positive

44
Q

Why is the process in electrochemical cells called a redox process?

A

Because oxidation and reduction happens simultaneously

45
Q

How do you make an electrochemical cell?

A

Can be made from 2 different metals dipped in a salt solution of their own ions and connected by a wire (external circuit)

46
Q

What are half-cells?

A

The 2 different sides of an electrochemical cell

47
Q

Why are the solutions connected by a salt bridge?

A

To allow ions to flow through and balance out the charges, thus completing the circuit

48
Q

In an electrochemical cell, what is the salt bridge made from?

A

A piece of filter paper soaked with KNO3 (aq)

49
Q

What is the cell potential?

A

The voltage between 2 half-cells in an electrochemical cell

50
Q

Describe the flow of electrons in an electrochemical cell

A

They flow from the more reactive metal to the less reactive metal

51
Q

Describe the properties of the electrode needed for an electrochemical cell

A

Needs to be a solid that conducts electricity

52
Q

How do you create a standard hydrogen electrode?

A

You can use non-metals as reactants in electrochemical cells. The gas can be bubbled over an inert electrode sitting in a solution of its aqueous ions

53
Q

What are the 5 steps to set up an electrochemical cell and then take measurements of voltage?

A

1: Get a strip of each of the metals you’re investigating
2: Clean the strips of metal using propanone
3: Place each electrode into a beaker filled with a solution of ions of the metal
4: Create a salt bridge to link the 2 beakers together by soaking a piece of filter paper with salt solution
5: Connect the electrodes to a voltmeter using wires and crocodile clips

54
Q

If you perform the experiment to take measurements of voltage from an electrochemical cell in standard conditions, what is the voltage equal to?

A

Standard cell potential for that cell

55
Q

Describe the reactions at the electrodes in an electrochemical cell

A

They are reversible reactions e.g. Zn2+ + 2e- ⇌ Zn

56
Q

What are electrode potentials?

A

The voltage measured when a half-cell is connected to a standard Hydrogen electrode

57
Q

Describe the electrode potential value for a metal electrode that is easily oxidised

A

Very negative electrode potential

58
Q

For a zinc/copper cell, which direction does the reaction go?

Electrode potential values:

Zn2+/Zn = -0.76V
Cu2+/Cu = 0.34V
A

The zinc electrode has a more negative electrode potential so it’s oxidised easier, so the reaction is backwards. Meaning Zinc is oxidised and the Copper is reduced

59
Q

Name a substance that could be used for the electrode in a half-cell involving solutions of 2 aqueous ions of the same element?

A

Platinum

60
Q

Why is Platinum suitable for use for the electrode in a half-cell involving solutions of 2 aqueous ions of the same element?

A

It’s a solid that conducts electricity and is inert

61
Q

What are half-cell reactions described as?

A

Reversible reactions

62
Q

What 3 factors affect the equilibrium position in a half-cell?

A

Temperature
Pressure
Concentration

63
Q

What changing the equilibrium in a half-cell affect?

A

The cell potential

64
Q

Why are standard conditions used for measuring electrode potentials?

A

So you always get the same value for the electrode potential because the equilibrium will be the same for each measurement, meaning you can compare different cells

65
Q

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

A

The voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode

66
Q

What is a standard hydrogen electrode?

A

An electrode where Hydrogen gas is bubbled through a solution of aqueous H+ ions under standard conditions

67
Q

What is the value given to the standard electrode potential of a hydrogen cell?

A

0V

68
Q

How can you use a standard hydrogen electrode to measure electrode potentials?

A

Because a hydrogen electrode has a potential of 0V, the voltage reading is equal to the standard electrode potential of whatever you’ve attached to the hydrogen half-cell

69
Q

What are the 3 conditions needed to measure electrode potentials under standard conditions?

A
  1. Any solutions of ions have concentration 1.00mol dm^-3
  2. Temperature = 298K (25°C)
  3. Pressure = 100kPa
70
Q

What 4 things are important when drawing a diagram to show how the standard electrode potential for a particular half-cell is measured?

A
  • Always put hydrogen electrode on left
  • Make sure to complete circuit (salt bridge, wires between electrode and voltmeter)
  • Label any solutions as being 1.00mol dm^-3
  • If half-cell contains aqueous ions of same element but different oxidation numbers, include a platinum electrode
71
Q

Why is a platinum electrode used in a standard hydrogen half-cell?

A

It is inert so is corrosion-resistant and catalyzes the proton reduction reaction

72
Q

When 2 half-equations are put together in an electrochemical cell, which one goes in the direction of oxidation (backwards)?

A

The half-equation with the most negative electrode potential

73
Q

When 2 half-equations are put together in an electrochemical cell, which one goes in the direction of reduction (forwards)?

A

The half-equation with the most positive electrode potential

74
Q

Describe the electrode potential of a more reactive metal

A

The more reactive a metal, the more it wants to lose electrons to form a positive ions, the more negative the standard electrode potential is

75
Q

Describe the electrode potential of a more reactive non-metal

A

The more reactive a non-metal, the more it wants to gain electrons to form a negative ion, so the more positive the standard electrode potential is

76
Q

What is the equation to work out the standard cell potential or e.m.f when 2 half-cells are connected together?

A

EΦcell = EΦreduced - EΦoxidised

EΦcell: Standard cell potential or e.m.f
EΦreduced: Standard electrode potential of half-cell which is reduced (has the more positive electrode potential)
EΦoxidised: Standard electrode potential of half-cell which is oxidised (has the more negative electrode potential)

77
Q

What is cell potential?

A

Voltage between 2 half-cells

78
Q

What are the 4 steps to predict whether a redox reaction will occur and to show which direction it’ll go?

A

1: Find the 2 half-equations and write them both as reduction reactions
2: Use electrochemical series to work out which half-equation has the more negative electrode potential
3: Write out half-equation of with more negative electrode potential as going in backwards direction and vice versa for more positive electrode potential
4: Combine both half-equations to form a redox equation

79
Q

Give 2 problems when predicting redox reactions causing you to potentially get the reaction wrong

A
  • The conditions are not standard

- The reaction kinetics are not favourable

80
Q

Why does not having standard conditions give a problem when predicting reactions?

A

It can cause the electrode potentials to change because the equilibrium position would be different

81
Q

Why does not having reaction kinetics that are favourable give a problem when predicting reactions?

A

The rate of reaction may be so slow that the reaction seems not to happen. Also, if a reaction has a high activation energy this would cause it to not happen

82
Q

What are energy storage cells?

A

A (rechargeable) battery that stores chemical energy to be converted into electricity

83
Q

How do you calculate the voltage of an energy storage cell?

A

Use same equation to work out the cell potential:

EΦcell = EΦreduced - EΦoxidised

84
Q

Why can some batteries be recharged?

A

Because the reactions that occur in them are reversible

85
Q

How do you recharge a battery?

A

A current is supplied to force electrons to flow in the opposite direction around the circuit and reverse the reactions. This is possible because none of the substances in a battery are used up or can escape

86
Q

How do you write the equations for recharging a battery?

A

Just reverse the equation for the storage cell

87
Q

Give 2 advantages of electrochemical cells

A

Cheap to make

Relatively high power densities

88
Q

Give a disadvantage of electrochemical cells

A

The production involves using toxic chemicals, which have to be disposed of when the battery has reached the end of its lifespan

89
Q

What is a fuel cell?

A

A cell that produces electricity by reacting a fuel with oxygen

90
Q

Give 3 advantages of using a fuel cell over a combustion engine

A
  • Fuel cells are more efficient at producing energy because energy is wasted during combustion as heat
  • Fuel cells produce a lot less pollution
  • For Hydrogen fuel cells, the only waste product is water
91
Q

What is the fuel in a hydrogen-oxygen fuel cell?

A

Hydrogen

92
Q

What are the 4 stages for how a Hydrogen fuel cell works?

A

1: At the anode, the platinum catalyst splits the H2 into protons and electrons
2: The polymer electrolyte membrane (PEM) only allows H+ across, forcing the electrons to travel around the circuit to get to the cathode
3: An electric current is created in the circuit from the flow of electrons
4: At the cathode, O2 combines with the H+ from the anode and the electron to form water, the only waste product

93
Q

What is the overall redox reaction for a hydrogen fuel cell in acidic or alkaline conditions?

A

1/2 O2 + H2 —> H2O

A Level Chemistry Year 2 (7 decks)

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