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Chemistry Physical > Electrochemical Cell > Flashcards

Electrochemical Cell Flashcards

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1
Q

Describe the set up of an electrochemical cell

A
  • 2 diff. metal ions connected by wire (external circuit) - e- flow from more reactive metal to less reactive
  • Dipped in salt solution of own ions connected by salt bridge (filter paper dipped in KNO<strong>3</strong> )
  • Voltmeter show v btw 2 half cells - EMF
  • Solutions of aqueous ions uses Pt electrode
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2
Q

Explain the function of the salt bridge

A
  • Allows ions to flow
  • Maintain charge balance
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3
Q

State a requirement of the soluble ionic compound used to make a salt bridge (1)

A

Must not react w/ electrolyte/ions in sol

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4
Q

Why is platinum used?

A

Inert- won’t react w/ ions

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5
Q

Which two reactions occur w/i an electrochemical cell?

A
  • Oxidation
  • Reduction
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6
Q

Are reactions that occur at each electrodes irreversible or reversible?

A

Reversible

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7
Q

How are electrochemical half-equations always written like?

A

Reduction reaction going in forward reaction

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8
Q

What happens when 2 half-equations are joined?

A
  • More -ve electrode potential (easily oxidised) is flipped
  • More +ve electrode potential (reduced) stays the same
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9
Q

What are electrode potentials measured against?

A

Standard hydrogen electrodes

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10
Q

Define standard electrode potential of a half cell

A

Voltage measured under standard conditions when the half-cell is connected to a standard hydrogen electrode

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11
Q

What are the standard conditions?

A
  • Conc: 1.00 mol dm-3
  • Temp: 298k (25oC)
  • Pressure: 100kPa
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12
Q

Best oxidising agent

A

Most +ve (left w/ charge)

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13
Q

Worst oxidising agent

A

Most -ve (left w/ charge)

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14
Q

Best reducing agent

A

Most -ve (right)

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15
Q

Worst reducing agent

A

Most +ve (right)

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16
Q

EMF

A

EMF = R - L

17
Q

Outline how to draw electrochemical cells

A

R | O || O | R

  • Double lines: salt bridge
  • Lines: state change
  • More -ve on left exp. w/ standard H
  • O = ions
18
Q

Most -ve

A
  • Anode
  • Most reactive
  • Oxidised
  • Loses e-
19
Q

Most +ve

A
  • Cathode
  • Least reactive
  • Reduced
  • Gains e-
20
Q

Sulfate ions: +0.17v Bromide ions: +1.07v

Justify why sulfate ions should not be capable of oxidising bromide ions (1)

A

E(Br2/Br-) > E(SO4-2/SO2)

21
Q

Outline the uses of lithium batteries

A
  • Rechargeable
  • Mobile phones, laptops + cars
22
Q

What are lithium batteries made up of?

A
  • Lithium cobalt oxide (LiCoO2) electrode + graphite electrode
  • Electrolyte: lithium salt in organic solvent
23
Q

What are the half equations for a lithium battery?

A
  • -ve electrode : Li → Li+ + e-
  • +ve electrode : Li+ + CoO2 + e- → Li+ [CoO2]-
24
Q

What happens when a battery is recharged?

A

Current force e- to flow in opposite direction round the circuit + reverse the reactions

25
Q

Why can’t non-rechargable batteries be recharged?

A

Reaction can’t be reversed

26
Q

Give a use of an alkaline hydrogen-oxygen fuel cell

A

Power electric vehicles

27
Q

Outline the structure of an alkaline hydrogen-oxygen fuel cell

A
  • H2 + O2 fed to 2 seperate Pt containing electrodes
  • Electrodes seperated by anion-exchange membrane that allows anions (OH-) + water to pass through it but not H + O gas
  • Electrolyte: aqueous alkaline solution (KOH)
  • Anions flow towards -ve electrode
28
Q

In an alkaline hydrogen-oxygen fuel cell which electrode is hydrogen fed to?

A

-ve electrode

29
Q

In an alkaline hydrogen-oxygen fuel cell which electrode is oxygen fed to?

A

+ve electrode

30
Q

What is the half equation for the reaction that occurs at the positive electrode for a hydrogen-oxygen fuel cell that operates in alkaline conditions?

A

O2 + 2H2O + 4e- → 4OH-

31
Q

What is the half equation for the reaction that occurs at the negative electrode for a hydrogen-oxygen fuel cell that operates in alkaline conditions?

A

2H2 + 4OH- → 4H2O + 4e-

32
Q

What is the overall equation for reaction that occurs at a hydrogen-oxygen fuel cell that operates in alkaline conditions?

A

2H2 + O2 → 2H2O

33
Q

What is the half equation for the reaction that occurs at the negative electrode for a hydrogen-oxygen fuel cell that operates in acidic conditions?

A

H2 → 2H+ + 2e-

34
Q

What is the half equation for the reaction that occurs at the positive electrode for a hydrogen-oxygen fuel cell that operates in acidic conditions?

A

O2 + 4H+ + 4e- → 2H2O

35
Q

How is a current generated in a hydrogen-oxygen fuel cell?

A
  • Hydrogen produces e-
  • Oxygen accepts e-
36
Q

What are the advantages of using fuel cells in cars than the internal combustion engine?

A
  • More efficient - greater proportion of energy available from hydrogen-oxygen reaction is converted into kinetic energy
  • Waste product is H2O - no toxic chem
  • Don’t need to be recharged
37
Q

What are the disadvantages of using a fuel cell?

A
  • Need energy to supply hydrogen + oxygen (electorlysis - uses electricity)
  • H is highly flammable/explosive
38
Q

Explain why a fuel cell doesn’t need to be recharged (1)

A

Hydrogen/fuel is continuously supplied

Chemistry Physical (11 decks)

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