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Chemistry - Module 5 > 23 - redox and electrode potentials > Flashcards

23 - redox and electrode potentials Flashcards

1
Q

what do oxidising agents do in terms of electrons

A

takes electrons from species being oxidised

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2
Q

what do reducing agents do in terms of electrons

A

adds electrons to the species being reduced

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3
Q

how to create an overall redox electrons from 2 half equations

A
  • balance electrons
  • add and cancel electrons
  • cancel any species on both sides of equation
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4
Q
A
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5
Q

how to form an overall redox equations from oxidation numbers

A
  • write initial equation
  • assign oxidation numbers
  • match the oxidation changes by balancing only that species
  • balance remaining atoms
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6
Q
A
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7
Q

what to take into account when predicting products

A
  • In aqueous redox reactions – H2O often formed
  • Balance O first then H+
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8
Q
A
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9
Q

what is reduced in the manganate (VII) redox titrations

A
  • MnO4 (-) reduced to Mn(2+)
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10
Q

write out the method for the manganate titration in full against a solution being analysed

A
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11
Q

why does manganate decolourise as it is being added

A

because it is reduced

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12
Q

colour change in mangante titration

A

purple to light pink

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13
Q

what occurs if insufficient amount of acid is added

A

solution will not be acidic enough and a brown solid, MnO₂, will be produced instead of Mn²⁺ ions.

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14
Q

why is sulfuric acid used instead of ethanoic acid

A

weak acid

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15
Q

Why can concentrated hydrochloric acid not be used instead of sulfuric acid

A

Concentrated hydrochloric acid cannot be used as the manganate (VII) ions will oxidise the chloride ions to chlorine.

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16
Q

determine percentage purity

A
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17
Q

determine value of x

A
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18
Q

in iodine-thiosulfate titrations what is oxidised and what is reduced

A

thiosulfate ions - oxidised

iodine - reduced

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19
Q

for an iodine thiosulfate titration - write out the half equations for oxidation and reduction and work out the overall equation

A
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20
Q

what can iodine-thiosulfate titrations be used for

A
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21
Q

how to find copper content in copper alloys - initial step

A

Alloy reacted + dissolved in nitric acid and neutralised to form Cu2+

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22
Q

method of iodine-thiosulfate titration against an oxidising agent

A
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23
Q

in an iodine-thiosulfate titration how to determine on an end point and method

A
  • iodine is being reduced back into iodide ions and brown colour faded to pale straw colour
  • add small amount of starch indicator
  • deep blue black colour formed due to presence of iodine
  • once all iodine has been reduced - blue black colour fades quickly
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24
Q

how does the iodine thiosulfate actually measure concentration of oxidising agent

A

measures how much thiosulfate needed to turn the solution from brown to colourless

oxidising agent reacts with iodide ions to form iodine (brown)

iodine (brown) is then reduced back into iodide by thiosulfate ions

when brown colour fades - it means the oxidising agent isnt producing any more iodine - because it is used up

so you know how much thiosufate exactly uses up the oxidising agent

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25
Q

how to analyse conc of household bleach experimentally

A
  • form solution of bleach
  • add potassium iodide
  • titrate this against sodium thiosulfate in burette - using previous method
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26
Q

what is bleach

A

oxidising agent

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27
Q

copper alloys

A

brass or bronze

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28
Q

how to find copper content of copper salts - initial step

A

Cu2+ made by dissolving copper salt in water OR insoluble copper compounds reacted with acids

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29
Q

write equations for the analysis of copper through an iodine-thiosulfate titration

A

oxising agent react with iodide

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30
Q
A
31
Q

how to find stoichiometry between two different equations

A

follow round like a circle - double then half etc

32
Q

what is a voltaic cell

A

converts chemical energy into electrical energy

33
Q

what do half cells contain

A

chemical species in redox half-equation

34
Q

why do in a full cell - the chemicals in the 2 half cells must be kept apart

A

 Or electrons flow in uncontrollable way + heat energy released instead of electrical energy

35
Q

draw zinc and copper half cells

A
36
Q

how to set up a Metal / metal ion half cell

A
  • Metal rod dipped into solution of its aq metal ion
37
Q

when you write a half cell in short form - you use a vertical line between the ion and the element - what is that

A

phase boundary between aqueous solution + metal

38
Q

when writing equation for half cell - what is the convention

A

ALWAYS WRITTEN THAT REDUCTION IS FORWARD REACTION

39
Q

what is a Ion / ion half-cells

A
  • Contains ions of same element in different oxidation states
     E.g. aq iron 2+ and aq iron 3+
40
Q

what electrode do we use for ion / ion half cell

A

platinum

41
Q

why do we use platinum

A
  • inert
  • no metal to transport electrons in / out of cell - need metal that is electrically conductive
42
Q

draw a half cell for Fe 3+ / Fe 2+ and write the half equation

A
43
Q

what is the standard electrode potential

A

Tendency to be reduced and gain electrons

OR

the e.m.f of half-cell connected to standard hydrogen half-cell under standard conditions

44
Q

what is THE standard half cell

A

half cell with hydrogen gas + solution with H+ ions

45
Q

what is the electrode potential on THE standard half cell

A

0V

46
Q

draw THE standard half cell

A
47
Q

standard conditions for standard electrode potential

A

Solution conc – 1moldm-3

Temp 298K

Pressure 100kPa / 1 bar

48
Q

how do we determine standard electrode potentials of half cells

A

e.m.f of half-cell connected to standard hydrogen half-cell under standard conditions

49
Q

if don’t know standard electrode potential values how do we know which species is oxidised + reduced

A
  • Electrode with more reactive metal loses electrons + is oxidised
    = Negative cathode
  • Electrode with less reactive metal gains electrons + is reduced
    = Positive anode
50
Q

how to measure standard electrode potentials - set up

A

Half-cell connected to standard hydrogen electrode

electrode connected by wire

solutions connected by salt bridge (strip of filter paper soaked in KNO3)

51
Q

purpose of wire

A

flow of electrons

52
Q

purpose of salt bridge

A

flow of ions

53
Q

what does a salt bridge contain and give example

A

contains concentrated solution of electrolyte that doesn’t react with any solution

E.g. strip of filter paper soaked in KNO3

54
Q

draw a set up to measure standard electrode potential of copper 2+

A
55
Q

the more negative standard electrode potential value…

A

Greater tendency to lose electrons + undergo oxidation

56
Q

the more positive standard electrode potential value…

A

lesser tendency to lose electrons + undergo oxidation

(more likely to be reduced)

57
Q

what standard electrode potentials do metals tend to have

A

– tend to have negative values + lose electrons

58
Q

what standard electrode potentials do non-metals tend to have

A

positive values + gain electrons

59
Q

write out method to measure standard cell potentials between two half cells

A
60
Q

from that method if the value is 1.1V what is that actually

A

E_cell

61
Q

formula for E_cell

A
62
Q

figure out what is feasible

A

C has a more positive electrode cell potential value - will be reduced

so that OXIDISING AGENT should react with the other REDUCING AGENTS

63
Q

how to write an overall redox equation for 2 half equations using electrode cell potential values

A
  • identify what’s being oxidised + what’s being reduced
  • rewrite half equations in that way for froward reaction = standard is reduction
  • combine two half equations
64
Q

limitations of these feasibility predictions

A
  • reaction rate
  • concentration
65
Q

how is reaction rate a limitation

A

 some reactions have large activation energy – very slow rate

66
Q

how is concentration a limitation

A

standard electrode potentials – 1moldm-3

if conc not that – value of electrode potential will be different from standard

if conc is bigger than 1 – shift to right + remove electrons + electrode potential is less negative

67
Q

whats another limitation

A
  • standard electrode potentials only apply to aqueous equilibria – not all of them are
68
Q

colour change in manganate titration

A

dark purple to pale pink

69
Q

what is formed when copper ion reacts with iodide

A

copper iodide and iodine

70
Q

what is copper iodide

A

white precipitate

71
Q

what colour is iodine

A

brown

72
Q

manganate colour change

A
73
Q

how do you know which is anode + which is cathode

A

Chemistry - Module 5 (7 decks)

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