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oxidation Flashcards

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1
Q

When does metal-metal displacement occur

A

If a solid metal is below the metal ion, then it will displace the metal ion into its solid metal form

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2
Q

When does metal-hydrogen displacement occur

A

If a solid metal is below the hydrogen ion, then it will displace the hydrogen ion into hydrogen gas

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3
Q

When does hydrogen-hallide displacement occur

A

If a halide ion is below the halogen, then it will displace the halogen

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4
Q

What is corrosion

A

The oxidation of a metal in the presence of air

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5
Q

Define an electrode

A

electrical conductors on which electron transfer occurs at the surface

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6
Q

Anode definition

A

electrode where oxidation occurs

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7
Q

What is the anode denoted as for galvanic and electrolytic cells

A

denoted negative in galvanic cells and positive in electrolytic cells

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8
Q

What is the cathode

A

electrode where reduction occurs

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9
Q

What is the cathode denoted as for galvanic and electrolytic cells

A

denoted positive in galvanic cells and negative in electrolytic cells

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10
Q

What is the electrolyte for galvanic cells

A

solution which contains ions that conduct charge

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11
Q

What is the electrolyte for electroyltic cells

A

molten salt or aqueous solution in which electrodes are immersed
- Allows anions to be attracted to the anode, cations attracted to the cathode

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12
Q

Definition of a salt bridge and 3 purposes of it

A
  • contains a non-reactive electrolyte solution
    • Maintains electrical neutrality
    • Allows anions to flow from the salt bridge to the anode and cations to flow from the salt bridge to the cathode
    • Completes the electrical circuit
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13
Q

External circuit definition

A

A conductive path allowing electron flow to occur between cells

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14
Q

Where do electrons go

A

Through the power supply to the cathode, being drawn from the anode

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15
Q

What electrode is used for solid metal and metal ion cells

A

metal electrode is used

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16
Q

What electrode is used for metal ions indifferent oxidation states

A

inert electrode used

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17
Q

What electrode is used for dissolved non-metal and its ions

A

inert electrode used

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18
Q

What electrode is used gaseous non-metal and its ions

A

inert electrode used

19
Q

What 2 factors does corrosion of ion require

A

oxygen + water

20
Q

What does water do in corrosion of water

A

Water acts as an electrolyte and salt bridge

21
Q

Oxidation and Reduction equations for Corrosion of iron in a normal medium

A
  • Oxidation: Fe(s) ⟶ Fe2+ + 2e-
  • Reduction: O₂ + 2H₂O + 4e- ⟶ 4OH-
22
Q

Precipitation reaction for Corrosion of iron in a normal medium

A

Fe2+ + 2OH- ⟶ Fe(OH)₂

23
Q

Further Oxidation reaction for Corrosion of iron in a normal medium

A

Fe(OH)₂ + OH- ⟶ Fe(OH)₃ + e-

24
Q

Overall Equation for Corrosion of iron in a normal medium

A

4Fe + 3O₂ + 2xH₂O ⟶ 2Fe₂O₃.xH₂O

25
Q

Oxidation and Reduction equations for Corrosion of iron in an acidic medium

A
  • Oxidation: Fe(s) ⟶ Fe2+ + 2e-
  • Reduction: O₂ + 4H+ + 4e- ⟶ 2H₂O (data sheet)
26
Q

5 Factors affecting Process of Rusting

A
  • Presence of oxygen: this is the oxidising agent - the greater the concentration of oxygen, the greater the rate of corrosion
  • Presence of water: increases the flow of ions between anodic and cathodic sites which increases the rate of corrosion
  • Presence of electrolytes: improves the efficiency of water as the salt bridge and increases the rate of corrosion
  • pH: the lower the pH, the greater the rate of corrosion as the reduction potential of water increases (i.e. its reduction becomes more favourable)
  • Presence of less or more reactive: if iron is in contact with a less reactive
27
Q

3 Factors which can prevent rusting

A
  • Inert, non-metallic coating - prevents contact with oxygen and water i.e. painting, plastic, oil
  • Inert, metallic coating - prevents contact with oxygen and water i.e. Cu, Sn, Pb
  • Galvanising - coating with more reactive metal, such as zinc - The zinc oxidises more readily and forms a protective layer over the iron
    Cathodic protection using a sacrificial anode
    Cathodic Protection using a DC current
28
Q

Explain Cathodic Protection using a sacrifical anode

A
  • a more reactive metal is placed in contact with the iron
    • The more reactive metal oxidises and must be replaced periodically
    • The iron now becomes the site for reduction and so O₂ is reduced at this site
    • A very damp or wet environment is needed to act as a salt bridge
29
Q

Explain Cathodic Protection using a DC current

A
  • the iron is connected to the negative terminal of a low voltage DC current
    • The excess of electrons prevents oxidation
    • The positive terminal of the power supply is attached to a piece of scrap metal or an inert material (Platinum) which causes the oxidation of H₂O, thus is oxidised over time
    • A very damp or wet environment is needed to act as a salt bridge
30
Q

In what scenario can inert, metallic coating be problematic

A

if the metallic barrier is damaged as it makes the point of contact with iron more anodic so iron will corrode faster

31
Q

Explain Downs Cell and anode + cathode reactions

A
  • Process of making sodium commercially available through the electrolysis of molten sodium chloride
    • Anode: Oxidation of Chlorine gas
    • Cathode: Reduction of Sodium
32
Q

Explain Electroplating and anode + cathode reactions

A
  • The process of electrolysing in order to place a thin film of metal on an object
    • Anode: Metal to be used as coating
    • Cathode: The object to be coated
33
Q

Explain Electrorefining and anode + cathode reactions

A
  • Involves purifying a metal where the impure metal anode is itself oxidised
    • Anode: Impure metal is oxidised where other metals fall to the bottom as anode slime
    • Cathode: Pure metal
34
Q

Why cant other metals be reduced in electrorefining

A

Other metals cannot be reduced as voltage is carefully controlled

35
Q

Explain Primary Cells and anode + cathode reactions

A
  • Non-rechargeable cells typically used as disposable batteries
    • Anode: Oxidation of zinc
    • Cathode: 2MnO₂ + 2NH₄ + 2e- —> Mn₂O₃ + 2NH₃ + H₂O
36
Q

2 reasons Why can’t primary cell battery be recharged

A
  • The Mn₂O₃ does not revert back to MnO₂ and instead produces MnOOH
  • Eventually ammonium chloride produces an acidic environment that corrodes the zinc casing
37
Q

What is a secondary cell

A

Rechargeable cells typically used as batteries

38
Q

Normal Use of a Secondary Cell equations (Anode/Cathode)

A
  • Anode: flip bottom lead equation
  • Cathode: Normal top lead equation
39
Q

Rechargeable Use of a Secondary Cell equations (Anode/Cathode)

A
  • Anode: flip top lead equation
  • Cathode: bottom lead equation
40
Q

Define Fuel Cells

A

Cells which do not store the oxidising or reducing agent and are more “environmentally friendly” in terms of energy production

41
Q

3 Examples of Galvanic and Electrolytic Cells

A

Galvanic Cell

Examples

  • Dry cell (Primary Cell)
  • Lead-acid accumulator (Secondary Cell)
  • Fuel cells

Electrolytic Cells

Examples

  • Electrorefining cells
  • Electroplating cells
  • Downs Cell
42
Q

3 Similarities between Galvanic and Electrolytic Cells

A
  • Oxidation occurs at the anode/reduction occurs at the cathode
  • Cations move towards the cathode/anions move towards the anode
  • Both have an External circuit in through which current flows
  • Electrolyte for transfer of ions
43
Q

3 Differences Between Galvanic + Electrolytic Cells

A

Galvanic cell reactions are spontaneous; Electrolytic cell reactions are not spontaneous

Galvanic cells generate a voltage/electric current; Electrolytic cells require an external voltage/electric current

Galvanic cells convert chemical energy to electrical energy; Electrolytic cells convert electrical energy to chemical energy

Charge of the anode is designated as negative; Charge on the anode is designated as positive

chem (5 decks)

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