Chapter 20Electrochemistry

20.1 Introduction to Electrochemistry

Electrochemistry• The branch of chemistry that deals with

electricity-related applications of oxidation-reduction reactions.

• Electrochemical Cells:A system of electrodes and electrolytes in which either chemical reactions produce energy or an electrical current produces chemical change

Half-Cell: a single

electrode immersed in a solution of its

ions

Components of Electrochemical Cells

Electrolyte Sol’n CuSO4

Electrolyte Sol’n

ZnSO4

Zn Electrode Anode- where oxidation takes

place

Cu ElectrodeCathode-

where reduction

takes place

Conducting Wire

Electrode: conductor used to establish electrical

contact with a nonmetallic part of the circuit.

Half-Cell: a single electrode immersed in a solution of its

ions

Cu ElectrodeCathode- written

as Cu+2/Cu

Zn Electrode Anode- written as

Zn+2/ZnOverall Cell Written as:

anode | cathodeZn | Cu

Chapter 20Electrochemistry

20.2 Voltaic Cells

Electrochemistry

Voltaic / Galvanic

Cell

Rxns that produce voltage spontaneously

Porous barrier which prevents the spontaneous mixing of the

aqueous solutions in each compartment, but allows the

movement of ions in both directions to maintain electrical

neutrality

• A chemical rxn that results in a voltage due to a transfer of electrons

• A chemical rxn that results in a voltage due to a transfer of electrons

Batteries • Two or more dry voltaic cells

• Zinc-Carbon Battery Zn → Zn+2 + 2e-

2MnO2 + H2O + 2e- → Mn2O3 + 2OH -

Batteries • Alkaline Battery- no carbon rod, smallerZn + 2OH - → Zn(OH)2 + 2e-

2MnO2 + H2O + 2e- → Mn2O3 + 2OH-

Batteries • Mercury Battery- no carbon rod, smallestZn + 2OH - → Zn(OH)2 + 2e-

HgO + H2O + 2e- → Hg + 2OH -

Fuel Cells• A voltaic cell where

reactants are constantly supplied and products are removed.

Rxns that turn chemical energy

into electrical energy

Cathode: O2 + 2H2O + 4e- → 4OH –

Anode: 2H2 + 4OH – → 4e- + 4H2O

Net: 2H2 + O2 → 2H2O

12

CorrosionFormation of Rust:

4Fe (s) + 3O2 (g) + xH2O → 2Fe2O3∙xH2O

Anode: Fe (s) → Fe+2 (aq) + 2e-

Cathode: O2 (g) + 2H2O (l) + 4e- → 4OH –

Prevention of CorrosionGalvanizing Process by which iron or any

metal is coated with zinc. Cathodic Protection

Since zinc is more easily

oxidized, it is a sacrificial anode.

Electrode Potentials• Reduction Potential: the tendency for the half-

reaction to occur as a reduction half-reaction in an electrochemical cell.

• Electrode Potential: the difference in potential between an electrode and its solution

• Potential Difference (Voltage): a measure of the energy required to move a certain electric charge between the electrodes, measured in volts.

• Standard Electrode Potential (E°): a half-cell measured relative to a potential of zero for the standard hydrogen electrode (SHE)

Standard Electrode Potential, E°

• Positive E° means hydrogen is more willing to give up its electron, so positive reduction potentials are favored. Naturally occurring rxns have a positive value.

E° cell

= E° cathode - E°

anode

• Negative E° means the metal electrode is more willing to give up its electron, this is not favored. These rxns prefer oxidation over reduction.

• When a half-cell is multiplied by a constant (for balancing) the E° value is NOT multiplied!

• When a rxn is reversed (flipped) the sign of the E° value switches.

• In a voltaic cell, the half-rxn with the more negative standard electrode potential is the anode, where oxidation occurs.

Standard Electrode Potential, E°Standard Electrode Potential, E°

Cell Potential• The potential voltage a rxn can produce.

Cu2+ + 2e- Cu Eo = .34 V

Ag+ + e- Ag Eo = .80V

Reduction potentials

Because this is a spontaneous process:

(Ag+ + e- Ag) x 2 Eo = .80V

Cu Cu2+ + 2e- Eo = -.34 V

Cu + 2Ag+ Cu2+ + 2Ag Eo = .46 V

Since both rxns are reduction, one

must be oxidation, flip it, positive

voltage must result from spontaneous

rxns

Cell Potential• The potential voltage a rxn can produce.Na+ + e- Na Eo = -2.71 V

Cl2 + 2e- 2Cl- Eo = 1.36 V

Because this is nonspontaneous process:

2Na+ + 2Cl- 2Na + Cl2 Eo = -4.07 V

(Na+ + e- Na) x 2 Eo = -2.71 V

2Cl- Cl2 + 2e- Eo = -1.36 V

Nonspontaneous, must end in

negative voltage. Flip one to become

oxidation. ** Fuel Cell!

Chapter 20Electrochemistry

20.3 Electrolytic Cells

Electrochemistry

Electrolytic Cell

Rxns that require an energy source to react

• When electric voltage is used to produce a redox reaction, it is called electrolysis

Batteries • Car Battery- rechargeable b/c the alternator reverses the ½ rxns and regenerates the reactants.

Discharge Cycle Rxn:

Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O

Electroplating• An electrolytic process in which a

metal ion is reduced and a solid metal is deposited on a surface

• Typically, an inactive metal is able to be ionized and then deposited on the surface of a more active metal to prevent corrosion.

Anode

Cathode

Silver ions are reduced at the cathode:

Ag+ + 1e- → Ag

Silver atoms are oxidized at the anode:

Ag → Ag + + 1e-

Voltaic vs. Electrolytic• If the positive battery terminal is attached to the cathode

of a voltaic cell, and the negative terminal is attached to the anode, the flow of electrons will change directions.

• Electrolytic cells need the electrodes attached to a battery, where voltaic is its own source of electrical power.

Voltaic = spontaneouschemical energy → electrical energy

Electrolytic = non-spontaneouselectrical energy → chemical energy

Electrolysis

Using a current to generate a redox reaction which otherwise would have a negative cell potential. i.e. electroplating & rechargeable batteries.

Anode: 6H2O → O2 + 4e- + 4H3O+

Cathode: 4H2O + 4e- → 2H2 + 4OH –