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chemistry - module 5 !! > redox & electrode potential °‧⭑.ᐟ > Flashcards

redox & electrode potential °‧⭑.ᐟ Flashcards

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1
Q

what are reducing agents in redox reactions?


A

electron donors

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2
Q

what are oxidising agents in redox reactions?


A

electron acceptors

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3
Q

what does oxidation involve in terms of electrons?


A

loss of electrons

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4
Q

what is the oxidation reaction of zinc?


A

Zn → Zn²⁺ + 2e⁻

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5
Q

what does reduction involve in terms of electrons?


A

gain of electrons

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6
Q

what is the reduction reaction of chlorine?


A

Cl₂ + 2e⁻ → 2Cl⁻

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7
Q

in the reaction Br₂ +2I⁻ → I₂ + 2Br⁻, which species is oxidised?


A

iodide ion

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8
Q

in the reaction Br₂ +2I⁻ → I₂ + 2Br⁻, which species is reduced?

A

bromine

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9
Q

how should oxidising and reducing agents be named?


A

full name of substance

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10
Q

what is the oxidising agent in the reaction involving Br₂ (aq) + 2e⁻ → 2Br⁻?

A

bromine

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11
Q

what is the reducing agent in the reaction involving 2I⁻ (aq) → I₂ (aq) + 2e⁻?

A

iodide ion

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12
Q

what is an oxidising agent?

A

species that causes another to oxidise

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13
Q

what is a reducing agent?

A

species that causes another to reduce

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14
Q

species that causes another to reduce

A

on the left

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15
Q

where are electrons placed in an oxidation half equation?

A

on the right

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16
Q

how do you balance half equations with varying oxygen amounts?

A

add OH⁻, H⁺ and H₂O

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17
Q

what is the half equation for MnO₄⁻ to Mn²⁺ in acidic conditions?

A

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + H₂O

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18
Q

what is the full balanced equation for the reduction of MnO₄⁻ and oxidation of C₂O₄²⁻?

A

2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 10CO₂ + 8H₂O

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19
Q

what is the half equation for the change of SO₄²⁻ to SO₂

A

SO₄²⁻ + 4H⁺ + 2e⁻ → SO₂ + 2H₂O

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20
Q

what is the common exercise in manganate redox titration?

A

redox titration between Fe²⁺ and MnO₄⁻

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21
Q

why is the manganate titration considered self-indicating?

A

significant colour change from purple to colourless

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22
Q

what is the role of acid in manganate titrations?

A

to supply 8H⁺ ions

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23
Q

why should only dilute sulfuric acid be used in manganate titrations?

A

other acids can cause inaccurate titration readings

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24
Q

what is the consequence of using nitric acid in manganate titrations?

A

oxidizes Fe²⁺ to Fe³⁺

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25
Q

how does the purple colour of manganate affect titration readings?

A

makes it hard to see the meniscus

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26
Q

what indicates the end point of the manganate titration?

A

first permanent pink colour appears

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27
Q

what is the common exercise in thiosulfate redox titration?

A

redox titration between I₂ and thiosulfate

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28
Q

what colour change occurs when starch is added near the end point of thiosulfate titration?

A

from blue/black to colourless

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29
Q

what are the steps in the manganate redox titration process?

A
  • add dilute sulfuric acid to the solution
  • titrate with MnO₄⁻ until the first permanent pink colour appears
  • record the volume of MnO₄⁻ used
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30
Q

what are the steps in the thiosulfate redox titration process?

A
  • add starch indicator near the end point
  • titrate with thiosulfate until the iodine fades
  • observe the colour change from blue/black to colourless
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31
Q

what is the balanced equation for the reaction in manganate titration?

A

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

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32
Q

what happens if insufficient sulfuric acid is used?

A
  • solution is not acidic enough
  • MnO₂ will be produced instead of Mn²⁺
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33
Q

what is the effect of brown MnO₂ in the titration?

A

masks the colour change

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34
Q

why can’t weak acids like ethanoic acid be used in manganate titrations?

A

cannot supply large amount of hydrogen ions needed

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35
Q

why is concentrated HCl unsuitable for manganate titrations?

A

Cl⁻ ions would be oxidisied to Cl₂ by MnO₄⁻

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36
Q

what is the balanced equation for the reaction between thiosulfate and iodine?

A

2S₂O₃²⁻ (aq) + I₂ (aq) → 2I⁻ (aq) + S₄O₆²⁻(aq)

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37
Q

what are the strengths and weaknesses of using different acids in manganate titrations?

A

strengths
- dilute sulfuric acid provides sufficient H⁺ ions

weaknesses
- other acids can lead to inaccurate readings
- concentrated HCl produces Cl⁻
- nitric acid oxidises Fe²⁺

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38
Q

why is heating the conical flask to 60°C beneficial in manganate titrations?

A

it speeds up the initial reaction

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39
Q

how do you calculate the percentage of iron by mass in a nail sample?

A

% mass = (mass of Fe / mass of nail) × 100

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40
Q

what is the role of a salt bridge in an electrochemical cell?


A

connects the half-cells and conducts charge

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41
Q

what is the typical composition of a salt bridge?

A

filter paper soaked in potassium nitrate

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42
Q

why is a high resistance voltmeter used in measuring electrode potential?

A

prevent current flow and measure maximum potential

43
Q

what happens to the voltage when current flows in an electrochemical cell?

A

voltage falls to zero as reactants are used up

44
Q

what is the potential difference measured in an electrochemical cell?

A

difference between two electrodes’ potentials

45
Q

what is the standard potential of the hydrogen electrode?

A

0 volts

46
Q

what is the overall reaction for manganate titration with hydrogen peroxide?

A

2MnO₄⁻ (aq) + 6H⁺ (aq) + 5H₂O₂ → 5O₂ + 2Mn²⁺ (aq) + 8H₂O

47
Q

how does the reaction between MnO₄⁻ and C₂O₄²⁻ proceed overall?

A

2MnO₄⁻ (aq) + 16H⁺ (aq) + 5C₂O₄²⁻ (aq) → 10CO₂ (g) + 2Mn²⁺ (aq) + 8H₂O (l)

48
Q

how do you calculate the moles used in a titration?

A

moles = concentration x volume

49
Q

what is the standard hydrogen electrode (SHE) assigned potential?

A

0 volts

50
Q

what is the equilibrium reaction for the hydrogen electrode?

A

H₂ (g) ⇌ 2H⁺ (aq) + 2e⁻

51
Q

why is a platinum wire used in the SHE?

A

acts as a catalyst
absorbs hydrogen gas

52
Q

what are the conditions for a standard hydrogen electrode?

A
  • hydrogen gas at 100 kPa
  • H⁺ concentration at 1 mol dm⁻³
  • temperature at 298 K
53
Q

what is a secondary standard in electrochemistry?

A

a standard calibrated against the SHE

54
Q

why are standard conditions necessary for measuring electrode potentials?

A

ensure consistent redox equilibrium positions

55
Q

what are the standard conditions for measuring electrode potentials?

A

1 mol dm⁻³ ions, 298 K and 100 kPa

56
Q

what is the standard electrode potential measured against?

A

standard hydrogen electrode

57
Q

how is the standard electrode potential quoted in data books?

A

half equations or cell notation

58
Q

what does a more negative half-cell potential indicate?

A

it will oxidise and go backwards

59
Q

how do you calculate the cell potential (Ecell) for a spontaneous reaction?

A

Ecell = E(red) – E(ox)

60
Q

what does a positive Ecell indicate about a reaction?

A

reaction is feasible

61
Q

what is the half-equation for zinc oxidation?

A

Zn (s) → Zn²⁺ (aq) + 2e⁻

62
Q

what happens to electrons at the positive electrode during a redox reaction?

A

electrons are absorbed (gained)

63
Q

what is the trend in electrode potentials for reducing and oxidising agents?

A
  • strong reducing agents: most negative potentials
  • strong oxidising agents: most positive potentials
64
Q

how do you determine the strongest oxidizing agent from a series of potentials?

A

look for the most positive potential

65
Q

how do you determine the strongest reducing agent from a series of potentials?

A

look for the most negative potential

66
Q

what is the half-equation for lithium reduction?

A

Li⁺ + e⁻ → Li

67
Q

what is the half-equation for fluorine reduction?

A

F₂ + 2e⁻ → 2F⁻

68
Q

what does an increasing positive potential indicate about a species?

A

has a greater tendency to reduce

69
Q

what does an increasing negative potential indicate about a species?

A

has a greater tendency to oxidise

70
Q

how do you calculate Ecell from two standard electrode potentials?

A

Ecell = E(red) – E(ox)

71
Q

what is the significance of E° values in redox reactions?

A

higher E° indicates stronger oxidising agents
lower E° indicates stronger reducing agents

72
Q

what is the difference in voltage behaviour between fuel cells and ordinary cells?

A

fuel cells maintain constant voltage over time

73
Q

what type of cells are nickel-cadmium cells?

A

rechargeable cells

74
Q

what is the overall reaction during the discharge of nickel-cadmium cells?

A

2NiO(OH) + Cd + 2H₂O → 2Ni(OH)₂ + Cd(OH)₂

75
Q

what characterizes non-rechargeable cells?


A

reactions within them are non-reversible

76
Q

what types of electrochemical cells exist?

A
  • non-rechargeable (irreversible) cells
  • rechargeable cells
  • fuel cells
77
Q

what is e.m.f?

A

a measure of how far from equilibrium the cell reaction lies

78
Q

what does it mean if the e.m.f is more positive?

A

more likely the reaction is to occur

79
Q

what would happen if the concentration of the reactants were to increase?

A

it would increase the e.m.f

80
Q

what would happen if the concentration of the reactants were to decrease?

A

it would decrease the e.m.f

81
Q

are most Ecells endothermic or exothermic?

A

exothermic

82
Q

what would the effect of a temperature rise have upon the Ecell?

A

decrease in Ecell

83
Q

if the reaction has a high activation energy will it occur?

A

no

84
Q

if the Ecell positive and it indicates a reaction might occur, what possibility may occur?

A

reaction will not occur or will occur so slowly that effectively it doesn’t happen

85
Q

why is an uncatalysed reaction very slow?

A
  • reaction needs a collision between two negative ions
  • repulsion between the ions is going to hinder this
  • high activation energy
86
Q

how can a substance act as a homogenous catalyst?

A
  • electrode potential must lie in between the electrode
    potentials of the two reactants
  • it can first reduce the
    reactant with the more positive electrode potential
  • then in the second step oxidise the reactant with the more
    negative electrode potential
87
Q

what do both individual stages in the catalysed mechanism involve?

A
  • collision between positive and negative ions
  • will have lower activation energies
88
Q

what does using the E values to find a catalyst only show?

A

catalysts is possible
does not guarantee that rate of reaction will increase

89
Q

what are electrochemical cells?

A

commercial source of electrical energy

90
Q

what are the three types of cells?

A
  • non-rechargeable
  • rechargeable
  • fuel cells
91
Q

when are cells non-rechargeable?

A

reactions that occur with in them are non-reversible

92
Q

what are nickel-cadmium cells used for?

A

power electrical equipment such as drills and shavers

93
Q

what type of cells are nickel-cadmium cells?

A

rechargeable

94
Q

what are fuel cells?

A

uses the energy from the reaction of a fuel with oxygen to create a voltage

95
Q

what is the hydrogen fuel cell overall reaction?

A

2H₂ + O₂ → 2H₂O

96
Q

what are the hydrogen fuel cell half equations in acidic conditions?

A

2e⁻ + 2H⁺ → H₂
4e⁻ + 4H⁺ +O₂ → 2H₂O

97
Q

what are the hydrogen fuel cell half equations?

A

4e⁻ + 4H₂O → 2H₂ + 4OH⁻
4e⁻ + 2H₂O + O₂ → 4OH⁻

98
Q

how can fuel cells will maintain a constant voltage over time?

A

continuously fed with fresh O₂ and H₂ so maintaining constant
concentration of reactants

99
Q

how do fuel cells differ from normal cells in terms of voltage?

A

ordinary cells’ voltage drops over time as the reactant concentrations drop

100
Q

why can’t standard conditions be used in a hydrogen fuel cell?

A

rate is too slow to produce an appreciable current

101
Q

why are higher temperatures used in hydrogen fuel cells?

A

increase rate but the reaction is exothermic
applying le chatelier would mean the emf falls
higher pressure can help counteract this

102
Q

advantages of fuel cells

A
  • less pollution and less CO2
    (pure hydrogen emits
    only water whilst hydrogen-rich fuels produce only
    small amounts of air pollutants and CO2)
  • greater efficiency
103
Q

disadvantages of hydrogen fuel cells

A
  • storing and transporting hydrogen, in terms of safety,
    feasibility of a pressurised liquid and a limited life cycle of a
    solid ‘adsorber’ or ‘absorber’
  • limited lifetime (requiring regular replacement and
    disposal) and high production costs
  • use of toxic chemicals in their production

chemistry - module 5 !! (1 decks)

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