Chapter 20 Electrochemistry 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

Components of Electrochemical Cells Cu Electrode Cathode- where reduction takes place Conducting Wire Electrolyte Sol’n ZnSO4 Electrolyte Sol’n CuSO4 Half-Cell: a single electrode immersed in a solution of its ions Electrode: conductor used to establish electrical contact with a nonmetallic part of the circuit. Zn Electrode Anode- where oxidation takes place

Half-Cell: a single electrode immersed in a solution of its ions Cu Electrode Cathode- written as Cu+2/Cu Half-Cell: a single electrode immersed in a solution of its ions Overall Cell Written as: anode | cathode Zn | Cu Zn Electrode Anode- written as Zn+2/Zn

Chapter 20 Electrochemistry 20.2 Voltaic Cells

Voltaic / Galvanic Cell Rxns that produce voltage spontaneously Electrochemistry 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 Voltaic / Galvanic Cell Rxns that produce voltage spontaneously 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, smaller Zn + 2OH - → Zn(OH)2 + 2e- 2MnO2 + H2O + 2e- → Mn2O3 + 2OH-

Batteries Mercury Battery- no carbon rod, smallest Zn + 2OH - → Zn(OH)2 + 2e- HgO + H2O + 2e- → Hg + 2OH -

Rxns that turn chemical energy into electrical energy Fuel Cells Cathode: O2 + 2H2O + 4e- → 4OH – Anode: 2H2 + 4OH – → 4e- + 4H2O Net: 2H2 + O2 → 2H2O A voltaic cell where reactants are constantly supplied and products are removed. Rxns that turn chemical energy into electrical energy

Corrosion Formation 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 Corrosion Galvanizing 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.

Standard Electrode Potential, E° 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.

Because this is a spontaneous process: Cell Potential The potential voltage a rxn can produce. Reduction potentials Cu2+ + 2e-  Cu Eo = .34 V Ag+ + e-  Ag Eo = .80V Since both rxns are reduction, one must be oxidation, flip it, positive voltage must result from spontaneous rxns 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

Because this is nonspontaneous process: Cell Potential The potential voltage a rxn can produce. Na+ + e-  Na Eo = -2.71 V Nonspontaneous, must end in negative voltage. Flip one to become oxidation. ** Fuel Cell! Cl2 + 2e-  2Cl- Eo = 1.36 V Because this is nonspontaneous process: (Na+ + e-  Na) x 2 Eo = -2.71 V 2Cl-  Cl2 + 2e- Eo = -1.36 V 2Na+ + 2Cl-  2Na + Cl2 Eo = -4.07 V

Chapter 20 Electrochemistry 20.3 Electrolytic Cells

Rxns that require an energy source to react Electrochemistry When electric voltage is used to produce a redox reaction, it is called electrolysis Electrolytic Cell Rxns that require an energy source to react

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 Silver ions are reduced at the cathode: Ag+ + 1e- → Ag Silver atoms are oxidized at the anode: Ag → Ag + + 1e- Cathode

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 = spontaneous chemical energy → electrical energy Electrolytic = non-spontaneous electrical energy → chemical energy

Electrolysis Anode: 6H2O → O2 + 4e- + 4H3O+ Cathode: 4H2O + 4e- → 2H2 + 4OH – Using a current to generate a redox reaction which otherwise would have a negative cell potential. i.e. electroplating & rechargeable batteries.