Electrochemistry
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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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Electrochemistry

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Electrochemistry

Electrochemistry


Electrochemical cells chemical cells

Electrochemical Cells/Chemical Cells

  • Also called voltaic or galvanic cells

  • A redox reaction produces electricity

  • Occurs spontaneously


Electrochemical cell

Electrochemical Cell


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


Electrochemistry

  • 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


Electrochemistry

  • 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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