Chapter 5: Galvanic and Fuel Cells Flashcards

1
Q

Electrochemical cell

A
  • Device that converts chemical energy into electrical energy (or vice versa)
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2
Q

Galvanic cell

A
  • Type of electrochemical cell in which chemical energy is converted into electrical energy
  • Reactions that occur in them are spontaneous and exothermic

NOTE: Galvanic cells are also known as voltaic cells.

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3
Q

Battery

A
  • Combination of several cells in series to obtain a higher potential difference or voltage
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4
Q

Reactions in each half cell

A
  • The anode (-) is where oxidation occcurs
  • The cathode (+) is where reduction occurs

TIP: AN OIL RIG CAT = anode, oxidation is loss & reduction is gain, cathode.

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5
Q

Salt bridge purpose

A
  • Often made from filter paper soaked in a relatively unreactive electrolyte (e.g. KNO₃)
  • Allow cells to produce electricity by allowing ions to move between the two half-cells (balances charges)
  • Cations (e.g K⁺) move toward the cathode and anions (e.g. NO₃-) move toward the anode
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6
Q

What occurs when there is no salt bridge?

A
  • The reaction would not proceed due to buildup of charge
  • One half cell would accumulate a negative charge and the other would accumulate a positive charge
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7
Q

Electrode materials

A
  • Redox pair with a metalgiven metal is used
    • E.g. Ag⁺(aq)/Ag(s): silver electrode
  • Redox pair without a metalplatinum or graphite is used
    • E.g. Fe³⁺(aq)/Fe²⁺(aq): platinum or graphite electrode
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8
Q

Strong vs weak reductants and oxidants

A
  • Strong reductants donate electrons more readily than weak reductants
    • They have weak conjugate oxidising agents
  • Strong oxidants accept electrons more readily than weak oxidants
    • They have weak conjugate reducing agents
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9
Q

Predicting cell reactions

A
  • Reductionhigher in the electrochemical series
    • Goes forward →
  • Oxidationlower in the electrochemical series
    • Reversed ←

TIP: ‘Clockwise’

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10
Q

Potential difference

A

= higher half-cell E⁰ (oxidant) – lower half-cell E⁰ (reductant)
- Measures the tendency to push electrons into the external circuit
- It is the electromotive force between two points in a circuit

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11
Q

Standard electrode potential (E⁰)

A
  • The voltage measured when a half-cell is connected to a standard hydrogen half-cell at standard conditions
  • Exists when one half-cell has a greater tendency to push electrons into the external circuit than the other cell

NOTE: It is also known as electromotive force (emf) or voltage.

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12
Q

Limitations of predictions using standard electrode potentials

A
  • The E⁰ values in the electrochemical series are only under standard conditions (values and order on the series vary under different conditions)
  • Does not provide information about the rate of reaction
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13
Q

Fuel cells

A
  • Type of galvanic cell that generate electricity from redox reactions
  • Chemical energy is converted directly into electrical energy
  • Reactants are not stored and must be continuously supplied
  • Two types: acidic and alkaline
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14
Q

Fuel cell design

A
  • Electrodes – must be conducting and porous to allow hydrogen and oxygen to come into contact with the ions in the electrolyte
    • Size determines the size of the current drawn from cell
  • Electrolyte – allows ions from anode and cathode to come into contact (does not allow transfer of electrons)
    • E.g. KOH
  • Catalysts – often coat electrodes to increase the reaction rate and current produced by the cell
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14
Q

Advantages and disadvantages of fuel cells

A
  • Advantages
    • Efficiency of 40-60% due to direct energy conversion
    • Reduced greenhouse gas emissions
    • Can use a variety of fuels
    • Generate electricity as long as fuel is supplied (conventional batteries need to be recharged/replaced)
  • Disadvantages
    • Require a constant fuel supply
    • Expensive
    • Storage and safety concerns
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