Electrochemi Electrochemi stry stry
Jan 11, 2016
ElectrochemistElectrochemistryry
Electrochemical CellsElectrochemical Cells
Electrons are transferred between the Electrons are transferred between the particles being oxidized and reducedparticles being oxidized and reduced
Two typesTwo types– Spontaneous = Voltaic (or Galvanic) CellSpontaneous = Voltaic (or Galvanic) Cell
Uses a spontaneous redox reaction to produced Uses a spontaneous redox reaction to produced electricityelectricity
– Non-spontaneous = Electrolytic CellNon-spontaneous = Electrolytic Cell Uses electricity to force a redox reaction to occurUses electricity to force a redox reaction to occur
Voltaic CellsVoltaic Cells
Also called galvanic cellsAlso called galvanic cells A redox reaction produces electricityA redox reaction produces electricity Occurs spontaneouslyOccurs spontaneously
Voltaic Cell
Half CellsHalf Cells
Each ½ of the redox reaction occurs in a Each ½ of the redox reaction occurs in a separate container separate container – One for oxidation and one for reductionOne for oxidation and one for reduction
They are connected by a salt bridgeThey are connected by a salt bridge– Salt Bridge: allows ions to flow between the Salt Bridge: allows ions to flow between the
two cellstwo cells
ElectrodesElectrodes
Metals which provide a surface for Metals which provide a surface for oxidation or reduction to occuroxidation or reduction to occur– SolidsSolids– Oxidation Number = 0Oxidation Number = 0– AnodeAnode– CathodeCathode
ANODEANODE– Oxidation occurs at the anodeOxidation occurs at the anode– Negative electrodeNegative electrode
CATHODECATHODE– Reduction occurs at the cathodeReduction occurs at the cathode– Positive electrodePositive electrode
Red Cat – An OxReduction at the
CathodeOxidation at the
Anode
Flow of ElectronsFlow of Electrons
The electrodes are connected by a wireThe electrodes are connected by a wire
Electrons flow from the anode to the Electrons flow from the anode to the cathode through the wirecathode through the wire
Why does the cell produce Why does the cell produce electricity? electricity?
There is a difference of electric potential There is a difference of electric potential between the two electrodesbetween the two electrodes– Electrons will flow between the two electrodes Electrons will flow between the two electrodes
until equilibrium is reacheduntil equilibrium is reached– At equilibrium the cell’s voltage would be zeroAt equilibrium the cell’s voltage would be zero
An Ox -oxidation takes place…electrons
are lost.
Red Cat -reduction takes place…electrons
are gained.
Zn Zn2+ + 2e- Cu2+ + 2e - Cu0
- +Electrons released here by
oxidation
Electrons needed here for
reductione-
e-
e-
e- e- e- e-
e-
e-
e-
e-
Zn + CuSO4 Cu + ZnSO4
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 CellsElectrolytic Cells
Also called electrolysisAlso called electrolysis An electric current is used to produce An electric current is used to produce
a chemical reactiona chemical reaction– An electric current is used to force a An electric current is used to force a
non-spontaneous reaction to occurnon-spontaneous reaction to occur
Oxidation occurs at the anodeOxidation occurs at the anode Reduction occurs at the cathodeReduction occurs at the cathode Electrons flow from anode to cathodeElectrons flow from anode to cathode The cathode is the negative The cathode is the negative
electrodeelectrode The anode is the positive The anode is the positive
electrodeelectrode– This is opposite of the chemical cell This is opposite of the chemical cell
because the external current causes the because the external current causes the polarities to switchpolarities 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
• 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
2 H2O 2 H2 + O2
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)