chapter 20 electrochemistry 20.1 introduction to electrochemistry
TRANSCRIPT
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
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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 –