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A-Level OCR Chemistry > Module 5: Chapter 23 (Redox and Electrode Potentials) > Flashcards

Module 5: Chapter 23 (Redox and Electrode Potentials) Flashcards

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1
Q

What is oxidation?

A

loss of electrons or increase in oxidation number

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2
Q

What is reduction?

A

loss of electrons or decrease in oxidation number

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3
Q

What is an oxidising agent?

A
  • takes electrons from the species that is being oxidised
  • contains the species that is reduced
  • for example: oxygen, dichromate, manganate (VII)
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4
Q

What is a reducing agent?

A
  • adds electrons to the species being oxidised
  • contains the species that is being oxidised
  • for example: magnesium, zinc, sulfur dioxide, hydrogen and Fe2+
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5
Q

How do you write a redox equation from half-equations?

A
  1. balance the electrons
  2. add the equations and cancel out the electrons
  3. cancel any species that is on both sides of the equation
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6
Q

Write a full equation:
H2O2 + 2e- -> 2OH-
Cr^3+ + 8OH- -> CrO4^2- + 4H2O + 3e-

A

3H2O2 + 2Cr^3+ + 10OH- -> 2CrO4^2- + 8H2O

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7
Q

What are the rules for oxidation numbers?

A
  1. Group one elements = +1
  2. group two elements = +2
  3. F = -1
  4. H = +1 (unless in a metal hydride = -1)
  5. O = -2
  6. Cl (and descending halides) = -1
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8
Q

How to write a redox equation from oxidation numbers?

A
  1. summarise any info provided into an equation
  2. assign oxidation numbers
  3. balance only the species that contain elements that have changed oxidation numbers
  4. balance any remaining atoms (H+ or H2O may have to be added to balance)
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9
Q

Balance:

S + HNO3 -> HSO4 + NO2 + H2O

A

S + 6HNO3 -> HSO4 + 6NO2 + 2H2O

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10
Q

Balance:

BrO3^- + Br- + H+ -> Br2 + H2O

A

BrO3^- + 5Br- + 6H+ -> 3Br2 + 3H2O

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11
Q

Balance (in acidic solution):

MnO4^- + SO3^2- -> SO4^2- + Mn^2+

A

6H+ + 2MnO4^- + 5SO3^2- -> 5SO4^2- + 2Mn^2+ + 3H2O

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12
Q

Balance (in acidic solution):

Cr2O7^2- + Sn^2+ -> Sn^4+ + Cr^3+

A

Cr2O7^2- + 3Sn^2+ + 14H+ -> 3Sn^4+ + 2Cr^3+ + 7H2O

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13
Q

What happens overall in a manganate (VII) titration?

A
  • carried out under acidic conditions

- MnO4- ions are reduced to Mn2+ and the other chemical is oxidised

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14
Q

What is the general method for manganate (VII) titrations?

A
  1. standard solution of potassium manganate (VII) fills the burette
  2. the other solution is pipetted into a conical flask, along with an excess of dilute sulfuric acid (acid provides the H+ ions)
  3. during the titration the manganate (VII) solution reacts and is decolourised. (MnO4- is pink/purple, Mn2+ is very pale/pink and often appears colourless)
  4. the end point is the first permanent pink colour (no indicator required)
  5. Repeat until concordant titres are obtained
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15
Q

What is the equation for the overall manganate equation?

A

5Fe2+ + MnO4- + 8H+ -> 5Fe3+ + Mn2+ + 4H2O
OR
5(COOH)2 + 2MnO4^- -> 16H+ -> 10CO2 + 10H+ Mn^2+ + 8H2O

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16
Q

Calculating the percentage purity of a product from manganate titration (e.g. FeSO4•7H2O)

A

(use a full balanced equation as reference)

  1. Calculate the amount of MnO4- that reacted using mean titre in n = c x V
  2. Determine the amount of Fe2+ that reacted using the balanced equation
  3. Find out unknown information (mol used in the titration scaling up to what has been used in the standard solution, then find the mass of the substance in the
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17
Q

Question:
A 0.1203g sample of (COOH)2•xH2O was dissolved in 25.0cm^3 of 1.0 moldm^-3 H2SO4 in a conical flask. The contents were heated to 60ºC and then titrated against 0.0200 moldm^-3 MnO4-. The mean titre was 19.10cm^3. Determine the value of x.
Half equations:
MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O and (COOH)2 -> 2CO2 + 2H+ + 2e-

A

(COOH)2•2H2O (x = 2)

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18
Q

What is the equation for the iodine/thiosulfate equation?

A

2S2O3^2- + I2 -> S4O6^2- + 2I-

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19
Q

What the solution colours in an iodine/thiosulfate equation?

A 
I2 = yellow-brown
I- = colourless
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20
Q

What are the 2 reducing agents manganate titrations can be used to analyse for?

A
  • Iron (II) iona, Fe^2+(aq)

- Ethanedioic acid, (COOH)2(aq)

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21
Q

What can the iodine/thiosulfate titration be used to determine?

A
  • the ClO- content in household bleach
  • the Cu2+ content in copper (II) compounds
  • the Cu content in copper alloys
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22
Q

What is the analysis of chlorate (I) ions?

A

ClO- + 2I- + 2H+ -> Cl- + I2 + H2O

  • ClO- is from NaClO (the active ingredient in bleach)
  • I2 produce reacts with S2O3^2- ions so therefore 1 mol of ClO- produces one mol of I2, which reacts with 2 moles of S2O3^2-
  • 1 mol ClO- = 2 moles of S2O3^2-
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23
Q

What is the analysis of copper (II) ions?

A

2Cu2+ + 4I- -> 2CuI + I2

  • CuI is a white ppt
  • I2 produced is titrated with thiosulfate
  • 1 mol of Cu2+ is equivalent to one mol of S2O3^2-
24
Q

What is a half-cell?

A

Half-cells contain all the chemical species present in a redox half-equation

25
Q

What is a voltaic cell?

A

These cells are made by connecting 2 different half-cells which then allows electrons to flow
- these cells convert chemical energy to electrical energy (therefore a type of electrochemical cell)

26
Q

What should chemicals in 2 half-cells be separated?

A

If they were not kept apart electrons would flow in an uncontrolled way and heat energy would be released instead of electrical energy

27
Q

What is electrical energy?

A

Electrical energy results from the movement of electrons and chemical energy is required to transfer electrons from one species to another

28
Q

What does the vertical like “|” in notations mean?

A

represents the phase boundary between the solid and the solution or ions for example

29
Q

What are metal/metal ion half-cells?

A
  • These half-cells consist of a metal rod dipped into a solution of its aqueous metal ions, e.g. Zn2+(aq) | Zn(s)
  • at the phase boundary, where the metal is in contact with its ions, an equilibrium is set up
  • by convention the forwards reaction shows reduction and the backwards reaction shows oxidation (Zn2+ + 2e- ⇌ Zn)
30
Q

What happens in an isolated half-cell?

A

There is no net transfer of electrons into or out of the metal

31
Q

What happens when 2 half-cells are connected?

A

Electrons are likely to flow (movement depends on relative tendency of each electrode to release electrons)

32
Q

What are ion/ion half-cells?

A
  • These contain ions of the same element in different oxidation states
  • There is no metal to transport electrons so an inert metal electrode is used (usually a platinum electrode)
  • | represents the phase boundary between the solution and the platinum electrode
33
Q

How is an electrochemical cell established?

A
  • connecting the 2 electrodes of the half-cells (e.g. Cu and Zn) by a wire via a high resistance volt meter (allows flow of electrons)
  • a salt bridge (normally KNO3) is used to complete the circuit
  • half-cells must be kept apart so that the flow of electrons can be controlled
34
Q

What happens with half-cells with gas in them?

A
  • a platinum electrode is used

- gas is introduced to the half-cell

35
Q

Describe and draw the standard hydrogen half-cell?

A
  • H2 gas introduced at 298K and at 1 atm through a glass tube with holes to let bubbles of H2 gas escape
  • platinum electrode
  • acid solution containing 1.0 moldm^-3 of H+(aq)

H+(aq), H2(g) | Pt(s) half-cell

36
Q

What is the redox equilibrium equation for the hydrogen half-cell?

A

2H+(aq) + 2e- ⇌ H2(g)

37
Q

How do you work out which electrode has the greatest tendency to gain/lose electrons?

A

In an operating cells

  • more reactive metal, loses electrons (negative electrode)
  • less reactive metal, gains electrons (positive electrode)
38
Q

How can the electrode potential, E, of a half-cell be found?

A
  • connecting it to a hydrogen-half cell under standard conditions
  • if the e- flow is from the half-cell to hydrogen: half cell has a negative standard electrode potential
  • if e- flows from hydrogen to half-cell: half-cell has a positive standard electrode potential
39
Q

What is the definition of standard electrode potential (E^θ)?

A

the emf (electromotive force, measured in V) of a half-cell compared with a standard hydrogen half-cell measured at 298K, with solution concentration of 1.0moldm^-3 and a gas pressure of 100KPa

40
Q

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

A

0.00V

41
Q

What can standard electrode potential values tell us?

A

more negative Eθ value = the greater tendency to lose electrons (and undergo oxidation) (negative electrode)
more positive Eθ value = the greater tendency to gain electrons (and undergo reduction) (positive electrode)

42
Q

What is the standard cell potential Eθcell?

A
  • cells are created when 2 half-cells are joined via a salt bridge and wires via a high resistance voltmeter
  • Eθcell is the voltage difference between the standard electrode potentials of the 2 half-cells
  • Eθcell is always positive
43
Q

Which way do electrons always flow?

A

always flow from the half-cell with the more negative standard electrode potential to the half-cell with the more positive standard electrode potential

44
Q

How is Eθcell calculated?

A

Eθcell = (more positive Eθ value) - (more negative Eθ)

45
Q

What carries charge in the salt bridge?

A

ions

46
Q

Key points from electrode potential data

A
  • represents the redox system in each half-cell
  • more negative Eθ = weaker oxidising agent
  • more positive Eθ = stronger oxidising agent
47
Q

How is reaction rate a limitation of predicting Eθ values?

A
  • a reaction between redox systems from predicted Eθ values will have a negative free energy ΔG and hence should be feasible
  • however a very high activation energy may cause a very slow reaction rate, as it it too small for it too noticeably occur at room temp
48
Q

How is concentration a limitation of predicting Eθ values?

A
  • standard conditions are not always satisfied
  • concentrations lower than 1.0 will shift equilibrium to the left, increasing the number of e- and making the electrode potential more negative
  • concentration above 1.0, equilibrium shifts to the right, decreasing the number of e- and making Eθ more positive
49
Q

What is another limitation of predicting Eθ values?

A
  • Eθ applies to aqueous equilibria

- many reactions take place not in aqueous solution

50
Q

What are primary cells?

A
  • non rechargeable
  • electrical energy is produced by oxidation and reduction at the electrodes
  • reactions can’t be reversed so chemicals wll be used up
  • batteries are either discarded or recycled after one use
51
Q

What are primary cells mostly used for?

A

low-current, low storage devices, e.g. wall clocks or smoke detectors

52
Q

What are primary cell mostly based on?

A
  • alkaline based on zinc and manganese dioxide (Zn/MnO2) and a potassium hydroxide alkaline electrolyte
  • the MnO2/Mn2O3 redox system is positive (reduced)
    2MnO2(s) + H2O(l) + 2e- -> Mn2O3(s) + 2OH-(aq)
  • the Zn/ZnO redox system is negative (oxidised)
    Zn(s) + 2OH-(aq) -> ZnO(s) +H2O(l) + 2e-
53
Q

What are secondary cells?

A
  • are rechargeable

- cell reactions producing electrical energy can be reversed during recharging as chemicals as regernrated

54
Q

What are common examples of secondary cells?

A
  1. Lead-acid batteries: car batteries
  2. Nickel-cadmium cells (NiCd) and Nickel-metal hydride (NiMH): cylindrical batteries used in radios, torches, etc
  3. Lithium-ion and lithium-ion polymer cells: used in modern appliances, laptops, tablets, cameras, mobile phones and being developed for cars
55
Q

What are fuel cells?

A
  • use energy from the reaction of a fuel with oxygen to generate a voltage
  • can be run continuously provided that the fuel and O2 are continually supplied to the cell
  • fuel cells do not have to be recharged
  • fuel and O2 flow into the cell, products flow out and electrolyte remains in the cell
  • hydrogen fuel is the most common fuel cell, no CO2 is produced
56
Q

What is the overall equation for alkali and acid hydrogen fuel cells?

A

H2(g) + 1/2 O2(g) -> H2O(l)

A-Level OCR Chemistry (21 decks)

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