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Electrochemistry. Electrochemical Cells/ Chemical Cells. Also called voltaic or galvanic cells A redox reaction produces electricity Occurs spontaneously. Electrochemical Cell. Half Cells. Each ½ of the redox reaction occurs in a separate container One for oxidation and one for reduction

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PowerPoint Slideshow about ' Electrochemistry' - torrance-prior


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Presentation Transcript
electrochemical cells chemical cells
Electrochemical Cells/Chemical Cells
  • Also called voltaic or galvanic cells
  • A redox reaction produces electricity
  • Occurs spontaneously
half cells
Half Cells
  • Each ½ of the redox reaction occurs in a separate container
    • One for oxidation and one for reduction
  • They are connected by a salt bridge
    • Salt Bridge: allows ions to flow between the two cells
electrodes
Electrodes
  • Metals which provide a surface for oxidation or reduction to occur
    • Solids
    • Oxidation Number = 0
    • Anode
    • Cathode
slide6
ANODE
    • Oxidation occurs at the anode
    • Negative electrode
  • CATHODE
    • Reduction occurs at the cathode
    • Positive electrode

Red Cat – An Ox

Reduction at the Cathode

Oxidation at the Anode

flow of electrons
Flow of Electrons
  • The electrodes are connected by a wire
  • Electrons flow from the anode to the cathode through the wire
why does the cell produce electricity
Why does the cell produce electricity?
  • There is a difference of electric potential between the two electrodes
    • Electrons will flow between the two electrodes until equilibrium is reached
    • At equilibrium the cell’s voltage would be zero
zn cuso 4 cu znso 4

Zn Zn2+ + 2e-

Electrons needed here for reduction

Electrons released here by oxidation

Zn + CuSO4 Cu + ZnSO4

Red Cat -reduction takes place…electrons are gained.

An Ox -oxidation takes place…electrons are lost.

Cu2+ + 2e - Cu0

e-

e-

e-

e-

e-

-

e-

+

e-

e-

e-

e-

e-

batteries
Batteries
  • Use a redox reaction which produces electricity spontaneously
  • Batteries are recharged by reversing the reaction
  • Dry Cell (Acid or Alkaline), Lead Storage (Car), Rechargeable (Ni/Cd)
corrosion
Corrosion
  • Oxidation of a metal
  • Metal combines with element (usually oxygen)

Example: 4Fe + O2 2Fe2O3 (rust)

prevention of rust
Prevention of Rust
  • Cover the metal – paint, oil, another (more reactive) metal
  • Cathodic Prevention
    • metal is placed in contact with a more reactive metal
    • That metal will be oxidized (acts as the anode), the original metal acts as the cathode
  • Alloys – mixture of metals
    • Brass, stainless steel (Fe + Cr), cast iron (C + Si)
electrolytic cells
Electrolytic Cells
  • Also called electrolysis
  • An electric current is used to produce a chemical reaction
    • An electric current is used to force a non-spontaneous reaction to occur
slide14
Oxidation occurs at the anode
  • Reduction occurs at the cathode
  • Electrons flow from anode to cathode
  • The cathode is the negative electrode
  • The anode is the positive electrode
    • This is opposite of the chemical cell because the external current causes the polarities to switch
electroplating
Electroplating
  • Object to be plated is the CATHODE, negative
  • Metal to be plated onto the object is the ANODE, positive
  • Solution must contain ions of the metal to be plated
silver plating
Silver Plating
  • Cathode =
  • Anode =
  • Solution =
  • What happens to the mass of each electrode during the reaction?
electrolysis of water
Electrolysis of Water

2 H2O  2 H2 + O2

  • The H+ is reduced at the (-) cathode, producing H2 (g), which is trapped in the tube
  • The O2 is oxidized at the (+) anode, yielding O2 (g), which is trapped in the tube
hydrogen fuel cells
Hydrogen Fuel Cells
  • Uses hydrogen gas as the fuel

2 H2 + O2  2 H2O

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