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A Level - Chemistry > 14: Redox 2 > Flashcards

14: Redox 2 Flashcards

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1
Q

What are the two options for balancing redox reactions?

A

Half reaction method

Change in oxidation method

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2
Q

What is the change in oxidation method for balancing redox equations?

A

Write out unbalanced reaction

Assigned oxidation numbers

Draw brackets and connect redox atoms

Find common factor, assign stoichiometry

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3
Q

What is the method for balancing equations in acidic conditions?

A

Write unbalanced equation

Assign oxidation numbers

Find common factor, assign stoichiometry

Add H2O to balance O atoms

Acidic solution balances H+ atoms from H2O

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4
Q

What is the method for balancing equations in basic conditions?

A

Write unbalanced equation

Assign oxidation numbers

Find common factor, assign stoichiometry

Add OH- to balance O atoms

H2O added to balance H and O

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5
Q

How is an equilibrium set up with metals?

A

Piece of metal is dipped into a solution of its metal ions

Metal forms positive ions and go into solution

Metal ions gain electrons and form metal

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6
Q

What is the half equation for a metal in a solution of metal ion?

A

Mn+ (aq) + ne- → M (s)

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7
Q

Where does the position of equilibrium effect the charge in a metal equilibrium?

A

Equilibrium to left, metal has -ve charge due to build up of electrons on the metal

Equilibrium to right, +ve charge builds up on metal as electrons are used up to form metal from metal ions

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8
Q

What is a negative and positive electrode potential in reference to equilibrium with a metal?

A

Eq. lies on left with -ve charge, -ve electrode potential

Eq. lies on right with +ve charge, +ve electrode potential

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9
Q

How does the position of equilibrium of a metal in solution show high reactivity?

A

More reactive metal tend to forms Mn+ ions

Negative charge builds up on the metal -ve electrode potentials

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10
Q

How does the position of equilibrium of a metal in solution show low reactivity?

A

Low reactive metal tend not to forms Mn+ ions

Positive charge builds up on the metal +ve electrode potentials

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11
Q

What is potential difference?

A

Difference between the positive and negative electrode

Measured in voltage

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12
Q

What are the three types of electrode?

A

Metal electrodes

Gas electrodes

Redox electrodes

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13
Q

What is a metal electrode?

A

Consists of metal surrounded by a solution of its ions

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14
Q

What is a gas electrode?

A 
Inert metal (platinum) is electrode to allow electrons to flow
Used for a gas and a solution of its ions
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15
Q

Which inert metal is used normally for the gas electrodes?

A

Platinum

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16
Q

What are redox electrodes?

A

For two ions of the same element where both are present in solution

Inert metal electrode allows flow of electrons

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17
Q

What is a half cell?

A

One of two electrodes in an electrochemical cell

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18
Q

What is reduction?

A

When a species gains electrons

Oxidation number decreases

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19
Q

What is oxidation?

A

When a species loses electrons

Oxidation number increases

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20
Q

How do s-block metals tend to react (terms of redox)?

A

Oxidised - lose electrons to form positive ions with charge same as group number

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21
Q

How do p-block elements tend to react (terms of redox)?

A

Metals react by losing electrons to form +ve ion

Non-metals react by gaining electrons to form -ve ion

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22
Q

How do d-block elements tend to react (terms of redox)?

A

Form ions with variable oxidation states

Tend to form +ve ions with +ve oxidation numbers

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23
Q

What reactions occur in a cell?

A

Always a reduction and oxidation reaction occurring

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24
Q

What charge is the anode in an electrochemical cell?

A

Negative electrode

Oxidation always occurs here

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25
Q

What charge is the cathode in an electrochemical cell?

A

Positive electrode

Reduction always occurs here

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26
Q

Where do electrons flow in a circuit?

A

From anode to cathode (more to least reactive metal)

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27
Q

Which electrode does the more reactive metal form?

A

Anode - as gives up its electrons more easily

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28
Q

Which electrode does the less reactive metal form?

A

Cathode - accepts electrons more easily

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29
Q

What is cell potential?

A

Voltage between the two half-cells Measure direction of flow of electrons

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30
Q

What is the convention of drawing half-cells?

A

Oxidation half-cell (anode) on the left Reduction half-cell (cathode) on the right

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31
Q

How are equations at electrodes presented?

A

Presented as reductions Reversible arrows show they can go in both directions

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32
Q

What are the steps to setting up an accurate electrochemical cell?

A

Strips of metals to investigate are cleaned using sandpaper

Clean grease/oil on electrodes with propanone, do not touch surface of metal with hands

Place into beaker with corresponding metal ions

Salt bridge made and electrodes connected to a voltmeter

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33
Q

What is a salt bridge?

A

Way to distribute ions to the different half-cells

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34
Q

What is the set up of a salt bridge?

A

Piece of filter paper soaked with unreactive ions Tube containing unreactive ions in an agar gel

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35
Q

Why is a salt bridge used?

A

Balance +ve and -ve charge in half cell solutions to maintain electrode potentials

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36
Q

Which compounds/ions are used in salt bridges?

A

KNO3 Forms unreactive ions

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37
Q

How are half-cell electrode potentials measured?

A

Measured relative to a particular half-cell (SHE)

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38
Q

What is SHE?

A

Standard hydrogen electrode Electrode chosen to be the primary standard which all others are measured against

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39
Q

Why are electrodes measured under standard conditions?

A

Various factors affect the electrode potential of a half-cell Due to half cells being in equilibrium

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40
Q

What are the standard conditions for a half-cell?

A

Cell conc - 1 M of ions involved in half-equation

Cell temp - 298K

Cell pressure - 100kPa (only affects half-cells with gases)

High resistance voltmeter

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41
Q

Why is a high resistance voltmeter used?

A

Must be done under zero-current conditions Allows measurement of full pd (emf), no current drawn from cell

42
Q

What are the standard conditions in a SHE?

A

1M solution of H+ atoms

298K

H2 at 100kPa

Salt bridge

Platinum electrode

43
Q

What is the half equation for the SHE?

A

2H+ (aq) + 2e- -> H2 (g)

44
Q

What are the conventions to drawing a cell?

A

|| separates half cells

Reduction is on the right

[] show substance flowing over inert metal

R|O||O|R

separates species/phases

45
Q

What is the convention for a hydrogen and magnesium electrochemical cell?

A

Pt[H2 (g)] | 2H+ (aq) || Mg2+ (aq) | Mg(s)

46
Q

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

A

Voltage measured under standard conditions when half-cell connected to SHE

47
Q

What is the electrode potential of a SHE?

A

0V

48
Q

What is E(cell)?

A

Cell potential of an electrochemical cell

49
Q

What is the equation for E(cell)?

A

E(cell) = E(reduction) - E(oxidation)

50
Q

Why is the E(cell) value always positive?

A

More negative E value is being subtracted from more positive E value

51
Q

What is the electrode potential of more reactive metals and why?

A

More negative standard electrode potential Loses electrons more easily

52
Q

What is the electrode potential of less reactive metals and why?

A

More positive standard electrode potential Gains electrons more easily

53
Q

When is a reaction thermodynamically feasible?

A

The overall E(cell) value is positive

54
Q

How can E values be used to predict disproportionation?

A

If E(cell) is positive for different redox reactions then disproportionation occurs

55
Q

Why might the E(cell) value predict something is not feasible but in practice it is?

A

Conditions are not standard Reaction kinetics aren’t favourable - slow rate of reaction so may not seem to occur, or high activation energy

56
Q

What is cell potential related to?

A

E(cell) is directly proportional to ΔS(total)

E(cell) is therefore directly proportional ln(K)

57
Q

How are pipettes used?

A

Measure only one volume of solution Fill pipette just above line and drop level carefully to the line

58
Q

How is a burette moved?

A

Burettes measure different volumes and let you add solutions drop by drop

59
Q

How are titrations used to calculate conc of acid/alkali?

A

Know volume of an alkali with unknown conc titrated with acid of known conc Volume of acid needed to neutralise acid used to calculate the conc of the alkali Or vice versa with acid/alkali

60
Q

What does an oxidising agent do?

A

Accepts electrons and is reduced

61
Q

What does a reducing agent do?

A

Donates electrons and is oxidised

62
Q

Why are d-block transition metals used as oxidising/reducing agents?

A

Good at changing oxidation numbers Readily give out/accept electrons

63
Q

How can the amount of reducing agent be calculated from an impure source?

A

Measure amount of impure product

Add dilute sulphuric acid to make up 250cm3

Titrate against oxidising agent such as MnO4- or C2O7 2-

Calculate moles

64
Q

What are the oxidation agents used which have a colour change in titration?

A

Potassium manganate KMnO4, goes from purple to colourless

Potassium Dichromate K2Cr2O7, goes from orange to green

65
Q

What are the ratios for iron and common oxidising agents?

A

Manganate - 5Fe2+ : MnO4-

Dichromate - 6Fe2+ : Cr2O7 2-

66
Q

What are the oxidation states of Vanadium?

A

+5 = VO3- or VO2+

+4 = VO2+

+3 = V3+

+2 = V2+

67
Q

What are the colours of the oxidation states of Vanadium?

A

Yellow = VO3- or VO2+ (oxidation state +5)

Blue = VO2+ (oxidation state +4)

Green = V3+ (oxidation state +3)

Violet = V2+ (oxidation state +2)

68
Q

What is the electronic configuration of Vanadium?

A

1s2 2s2 2p6 3s2 3p6 3d3 4s2

69
Q

What is the half equation for the changing on oxidation state of vanadium between +5 and +4?

A

VO2+ + 2H+ + e- -> VO2+ + H2O

70
Q

What is the half equation for the changing on oxidation state of vanadium between +4 and +3?

A

VO2+ + 2H+ + e- -> V3+ + H2O

71
Q

What is the half equation for the changing on oxidation state of vanadium between +3 and +2?

A

V^3+ + e- -> V^2+

72
Q

How is thiosulphate used for a titration?

A

XS iodide ions react with oxidising agent, liberating iodine

Iodine titrated against a standard solution of sodium thiosulphate

73
Q

What is the equation linking iodine and thiosulphate?

A

I2 + 2S2O32- -> 2I- + S4O62-

74
Q

What is the method for sodium thiosulphate titration?

A

Known volume of oxidising agent pipetted into a conical flask

Similar vol of dilute sulphuric acid added, with XS solid KI, swirled to ensure all oxidising agent reacts

Burette filled with standard solution of sodium thiosulphate

Thiosulphate added until brown fades to straw

Few drops of starch added to make it blue-black

Thiosulphate added until decolourised

75
Q

What colour change is seen when using thiosulphate?

A

Brown to pale straw before starch

Black-blue to colourless after starch

76
Q

What occurs when starch is added to an iodine solution?

A

Iodine forms complex with starch Forms a blue-black colour

77
Q

Why is manganate used instead of dichromate for some titrations?

A

Clearer colour change

78
Q

What is the method for titration with potassium manganate?

A

Known volume (usually 25cm3) of reducing agent is pipetted into a conical flask

Approx 25cm3 of dilute sulphuric acid is added

Burette filed with standard solution of potassium manganate

Manganate (purple) added to reducing agent bit by bit where it turns colourless

End-point when a permanent pink colour is seen

79
Q

What is the reaction which occurs in a potassium manganate titration?

A

Fe2+ reacts with MnO4 - ions turning it colourless

End-point when no Fe2+ is left in the solution

80
Q

What are the main types of cells?

A

Non-rechargeable

Rechargeable

Fuel cells

81
Q

What is a battery?

A

More than one cell which is joined together

82
Q

Why are non-rechargeable cells not used more than once?

A

Chemicals used up over time and emf drops

Once one or more chemicals run out it cannot be used

83
Q

What are the common chemicals found in non-rechargeable cells?

A

Zinc-carbon cell - standard and cheap with shorter life

Alkaline cell - higher cost but longer life

84
Q

How do rechargeable cells work?

A

Reactions are reversible

Reversed by applying an external current and regenerating the chemicals

85
Q

What are common types of rechargeable cells?

A

Lead-acid

Nickel-cadmium

86
Q

What is a fuel cell?

A

Cell which have a continuous supply of the chemicals into the cell

Chemicals are stored separately outside the cell

87
Q

What do the fuel cells not require like conventional cells?

A

Does not run out of chemicals

Not need recharging

88
Q

What is the most common fuel cell?

A

Hydrogen-oxygen fuel cell

89
Q

What is the set up of a hydrogen-oxygen fuel cell?

A

Hydrogen and oxygen gases fed into two separate platinum-containing electrodes

Electrodes separated by an anion-exchange membrane allowing only ions to move through it

90
Q

What are the two types of hydrogen-oxygen fuel cells?

A

Alkaline or acidic

91
Q

What are the half equations in an alkaline hydrogen-oxygen fuel cell?

A

H2 + 2OH- -> 2H2O + 2e-

O2 + 2H2O + 4e- -> 4OH-

Overall: 2H2 + O2 -> 2H2O

92
Q

What are the half equations in an acidic hydrogen-oxygen fuel cell?

A

H2 -> 2H+ + 2e-

O2 + 4H+ + 4e- -> 2H2O

Overall: 2H2 + O2 -> 2H2O

93
Q

What is the cell emf of a hydrogen-oxygen fuel cell?

A

Acidic and alkali: +1.23 V

94
Q

Which is the anion-exchange membrane and polymer electrolyte membrane?

A

Surface which allows passage of H+ or OH- ions to flow between electrodes in hydrogen-oxygen fuel cells

95
Q

Which electrode is hydrogen and oxygen delivered to in a hydrogen-oxygen fuel cell?

A

Hydrogen delivered to the anode Oxygen delivered to the cathode

96
Q

How is electricity generated in an acidic hydrogen-oxygen fuel cell?

A

H2 forms H+ ions at anode, donating electrons to it

Electrons move to cathode, creating a pd H+ ions move between electrodes through electrolyte to cathode

Electrons at cathode used to combine H+ and O2 to form water

97
Q

How is electricity generated in an alkali hydrogen-oxygen fuel cell?

A

O2 at cathode reacts with water and uses electrons from circuit to form OH-

OH- pass through anion-exchange membrane to negative anode

OH- reacts with H2 to form H2O and produces e- which go into the circuit and produces a pd

98
Q

What are the advantages/disadvantages to using cells?

A

Portable source of electrical energy

Waste issues

99
Q

What are the advantages/disadvantages to using non-rechargeable cells?

A

Cheap

Waste issues

100
Q

What are the advantages/disadvantages to using re-chargeable cells?

A

Less waste, cheaper in long run, lower environmental impact

Some waste issues at end

101
Q

What are the advantages/disadvantages to using hydrogen fuel cells?

A

Only waste product is water, no recharging, very efficient

Need constant supply of fuels, hydrogen is flammable and explosive and made using fossil fuels, high cost of fuel cells

102
Q

Define standard electrode potential

A

Emf of a half-cell measured relative to SHE All standard conditions (1M solution, 100kPa gas, 298K)

A Level - Chemistry (26 decks)

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