Transcript Electrochemistry

Electrochemistry
Electrochemical Cells
 Electrons are transferred between the
particles being oxidized and reduced
 Two types
– Spontaneous = Voltaic (or Galvanic) Cell
 Uses a spontaneous redox reaction to produced
electricity
– Non-spontaneous = Electrolytic Cell
 Uses electricity to force a redox reaction to occur
Voltaic Cells
 Also called galvanic cells
 A redox reaction produces electricity
 Occurs spontaneously
Voltaic Cell
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
Metals which provide a surface for
oxidation or reduction to occur
– Solids
– Oxidation Number = 0
– Anode
– Cathode
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
The electrodes are connected by a wire
Electrons flow from the anode to the
cathode through the wire
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 + CuSO4
Cu + ZnSO4
Red Cat -reduction
takes place…electrons
are gained.
An Ox -oxidation
takes place…electrons
are lost.
Zn
Zn2+ + 2e-
e-
e-
e-
Cu2+ + 2e
e-
-
Cu0
e-
-
ee-
+
eElectrons
released
here by
oxidation
e-
ee-
Electrons
needed
here for
reduction
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)
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
 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



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
•
•
•
•
Cathode =
Anode =
Solution =
What happens to the
mass of each
electrode during the
reaction?
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
• Uses hydrogen gas as the fuel
• Spontaneous (Electrochemical Cell)
2 H2(g) + O2(g)  2 H2O(g) + energy
Corrosion
• Oxidation of a metal
• Metal combines with element (usually
oxygen)
Example: 4Fe + O2
2Fe2O3 (rust)
Prevention of Corrosion
• 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)