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)