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redox and electrode potentials Flashcards

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1
Q

define oxidising agent

A

an oxidising agent accepts electrons and is reduced

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2
Q

define reducing agent

A

a reducing agent donates electrons and gets oxidised

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3
Q

define oxidation

A

oxidation is the loss of electrons

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4
Q

define reduction

A

reduction is a gain of electrons

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5
Q

what happens in a redox reaction?

A
  • electrons are transferred from one species to another
  • one element is oxidised whilst the other is reduced
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6
Q

write a half equation for the reaction of zinc metal with silver nitrate solution

A
  • first write full equation (if not given) and balance it:
    Zn(s) + 2AgNO3(aq) → Zn(NO3)2(aq) + 2Ag(s)
  • determine the redox numbers of zinc and silverzinc: 0 → 2+ (is oxidised)silver: 1+ → 0 (is reduced)
  • write half equations for each redox reaction
    Zn(s) → Zn2+(aq) +2e-
    2Ag+(aq) + 2e- → 2Ag (s)
  • balance the electron numbers in the reaction (if not balanced) and then cancel them out
    Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)
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7
Q

what can you add to balance a half equation

A
  • water (H2O)
  • electrons (e-)
  • H+ ions
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8
Q

write half equations for when acidified manganate (VII) ions (MnO4-) are reduced to Mn2+ by Fe2+ ions

A
  • Fe2+(aq) → Fe3+ (aq) + e-
  • MnO4-(s) → Mn2+ (aq)
  • balance oxygen: MnO4-(aq) → Mn2+(aq) + 4H2O(l)
  • balance hydrogen: MnO4-(aq) + 8H+(aq) → Mn2+(aq) + 4H2O(l)
  • balance charge: MnO4-(aq) + 8H+(aq) + 5e- → Mn2+(aq) + 4H2O(l)
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9
Q

how do you work out the concentration of a reducing agent

A

titrate a known volume of a reducing agent against a known concentration

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10
Q

describe a method that allows you to find out how many magnate ions are needed to react with a reducing agent

A
  1. measure a volume of the reducing agent using a pipette and place in a conical flask
  2. add dilute sulphuric acid to the flask
  3. gradually add MnO4- (oxidising agent) to the reducing agent using a burette and swirling the conical flask
  4. stop when the mixture in the flask slightly changes colour and record the volume of the oxidising agent added
  5. run a few titrations and calculate the mean volume of MnO4-
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11
Q

describe and explain the colour change in the titration using acidified potassium manganate(VII) solution

A

purple to colourless

  • MnO4- is purple
  • [Mn(H2O)6]2+ is colourless
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12
Q

describe and explain the colour change in the titration using acidified potassium dichromate(VI) solution

A

orange → green

  • Cr2O7 2- is orange
  • [Cr(H2O)6]3+ is violet/green
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13
Q

how can you make the colour change easier to spot in a titration

A

use a white tile underneath the conical flask

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14
Q

what is the iodine-sodium thiosulfate titration used for?

A

finding the concentration of an oxidising agent

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15
Q

the more concentrated an oxidising agent is the more….

A

the more concentrated an oxidising agent is the more ions will be oxidised by a certain volume of it

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16
Q

describe the method to oxidise iodide

A
  1. measure out a volume of potassium iodate (V) solution (KIO3) - the oxidising agent
  2. add this to an excess of acidified potassium iodide solution (KI)
  3. the iodate ions in the potassium iodate solution oxidise some of the iodide ions to iodine
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17
Q

what is the formula of sodium thiosulfate

A

Na2S2O3

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18
Q

what is the formula for thiosulfate ions

A

S2O3 2-

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19
Q

write the equation for the oxidation of iodide ions to iodine

A

IO3-(aq) + 5I-(aq) + 6H+(aq) → 3I2(aq) + 3H2O(l)

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20
Q

describe a method to titrate iodine with sodium thiosulfate

A
  1. from a burette add sodium thosulfate solution to the flask containing the oxidised iodide ions drop by drop
  2. when iodine colour fades to a pale yellow add 2cm of starch solution
  3. solution of conical flask will go dark blue to show iodine is still present
  4. add sodium thiosulphate one drop at a time until blue colour disappears (all iodine has been reacted)
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21
Q

write the equation showing the reaction between iodine with thiosulfate ions

A

I2(aq) + 2SO3 2-(aq) → 2I-(aq) + S4O6 2-(aq)

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22
Q

define a half cell

A

one half of an electrochemical cell constructed of a metal dipped in it’s ions, or a platinum electrode with 2 aqueous ions

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23
Q

when using a half cell with 2 aqueous ions what must you use?

A

when using a half cell with 2 aqueous ions you must use an inert but electrically conductive electrode

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24
Q

what happens when a rod of metal is dipped into a solution of its own ions?

A

an equilibrium is set up between the solid metal and the aqueous metal ions

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25
give an example of an inert but electrically conductive electrode
platinum
26
how are electrochemical cells created
they are created by joining two different half cells together by a wire, voltmeter and a salt bridge
27
what is a voltmeter used for
to measure the voltage between two half cells
28
what is the EMF / Ecell
the voltage between two half cells
29
describe the electron flow in an electrochemical cell
electrons flow from a more reactive metal to a less reactive one
30
write a half equation for zinc (s) to zinc (II)
Zn(s) ⇌ Zn2+ + 2e-
31
write a half equation for copper (II) to copper (III)
Cu2+(aq) ⇌ Cu3+(aq) + e-
32
what happens to an electrode that is being oxidised
- electrode becomes thinner - more ions are produced to make the electrons
33
what happens to an electrode that is being reduced
- electrode will get thicker - metal ions receive more electrons to turn into the metal solid
34
define standard electrode potential (E°)
the **voltage** of a **half cell** compared with a **standard hydrogen half cell** measured at **298K** with solution **concentration of 1moldm-3** and a **gas pressure of 100kPa**
35
what does the standard electrode potential tell us (E°)
how easily the half cell gives up electrons (is oxidised)
36
if the E° is more negative, what does it mean?
it is a better reducing agent (easier to oxidise)
37
if the E° is more positive, what does it mean?
it is a better **oxidising agent** - **GETS REDUCED**
38
what is the standard hydrogen electrode used for?
used as a reference to measure standard electrode potentials (E°)
39
what is a standard hydrogen electrode made of?
1. hydrochloric acid 1 mol dm-3 2. hydrogen gas at 100kPa 3. inert platinum electrode
40
why is a hydrogen half cell used as a standard half cell
- easy to control - pure - easy to reproduce
41
what is the E° of a standard hydrogen electrode
0.00V
42
how do you make a simple salt bridge?
soak a piece of filter paper in an aqueous solution of KNO3 or NH4NO3
43
explain why are salt brides necessary
- to complete the circuit by connecting the two solutions - enables charge to be transferred between the half cells - do not react with electrodes
44
why might other standard electrode electrodes be used instead of the standard hydrogen electrode
they are cheaper, easier, quicker to use and can provide just as good a reference
45
how do you calculate the voltage (emf) of a cell from E° values?
E° cell = E° positive - E° negative
46
when do you use a platinum electrode?
when both the oxidised and reduced forms of the metal are in aqueous solution
47
how do you predict if a reaction has occurred?
- take 2 half equations and find the species being reduced - calculate the E° value minus the E° value of the species that is being oxidised - if E° is bigger than 0.4V a reaction will occur
48
what are the 3 main types of electrochemical cells?
1. non-rechargable cells 2. rechargebale cells 3. fuel cells
49
describe how non rechargeable cells work
they provide electrical energy until all the chemicals have reacted
50
describe how rechargeable cells work
- chemicals in the cell provide electrical energy - when recharging the reactions of the cell become reversed
51
give examples of rechargeable cells
1. nickel and cadmium batteries 2. lithium ion batteries 3. lithium polymer batteries
52
explain why lithium is used in laptop batteries
- lithium has low density - so electrode is light and very reactive
53
what are the disadvantages of using lithium batteries
1. toxic if ingested 2. rapid discharge of current can cause fire/explosions
54
describe how fuel cells work
- cell uses external supplies of fuel and an oxidant - external supplies need to be continuously supplied
55
modern fuel cells are based on what type of fuels?
- hydrogen - hydrogen rich fuels
56
what are the reactions that take place at the two electrodes in an alkaline hydrogen fuel cell?
2H2 + 4OH- → 4H2O + 4e- O2 + 2H2O + 4e- → 4OH-
57
what are the disadvantages of fuel cells
1. hydrogen is a **flammable gas** with a low boiling point so it is **hard and dangerous to store and transport** 2. hydrogen is **expensive** to buy 3. fuel cells have a **limited lifetime** and use **toxic chemicals in their manufacture**
58
what is the reason that some cells cannot be recharged?
- reaction of the cell is not reversible - a product is produced that either dissipates or cannot be converted back into the reactants
59
why might the e.m.f (voltage) of a cell change after a period of time?
concentrations of the ions change - the reagents are used up
60
how can the e.m.f (voltage) of a cell be kept constant
reagents are supplied constantly so the concentration of the ions are constant
61
what are the advantages of electrochemical cells
1. they are **more efficient** at producing energy than conventional combustion engines because **less energy is wasted during combustion as heat** 2. they produce **less pollution (CO2)**
62
explain the advantages of using a platinum electrode
1. its **inert** so doesn’t affect the **redox reaction** 2. **conducts electricity so provides a way to add/remove electrons** 3. electrode is coated in **platinum black** which is **porous**, provides a **large surface area for the redox reaction**

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