Presentation on theme: "Chapter 20 Electrochemistry"— Presentation transcript:
1 Chapter 20 Electrochemistry 20.1 Introduction to Electrochemistry
2 ElectrochemistryThe 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
3 Components of Electrochemical Cells Cu ElectrodeCathode- where reduction takes placeConducting WireElectrolyte Sol’n ZnSO4Electrolyte Sol’n CuSO4Half-Cell: a single electrode immersed in a solution of its ionsElectrode: conductor used to establish electrical contact with a nonmetallic part of the circuit.Zn ElectrodeAnode- where oxidation takes place
4 Half-Cell: a single electrode immersed in a solution of its ions Cu ElectrodeCathode- written asCu+2/CuHalf-Cell: a single electrode immersed in a solution of its ionsOverall Cell Written as:anode | cathodeZn | CuZn ElectrodeAnode- written asZn+2/Zn
7 Voltaic / Galvanic Cell Rxns that produce voltage spontaneously ElectrochemistryPorous 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 neutralityVoltaic / Galvanic CellRxns that produce voltage spontaneouslyA chemical rxn that results in a voltage due to a transfer of electrons
8 Batteries Two or more dry voltaic cells Zinc-Carbon Battery Zn → Zn+2 + 2e-2MnO2 + H2O + 2e- → Mn2O3 + 2OH -
11 Rxns that turn chemical energy into electrical energy Fuel CellsCathode: O2 + 2H2O + 4e- → 4OH –Anode: 2H2 + 4OH – → 4e- + 4H2ONet: 2H2 + O2 → 2H2OA voltaic cell where reactants are constantly supplied and products are removed.Rxns that turn chemical energy into electrical energy
13 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 –
14 Prevention of Corrosion GalvanizingProcess by which iron or any metal is coated with zinc. Cathodic ProtectionSince zinc is more easily oxidized, it is a sacrificial anode.
15 Electrode PotentialsReduction 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 solutionPotential 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)
16 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° anodeNegative E° means the metal electrode is more willing to give up its electron, this is not favored. These rxns prefer oxidation over reduction.
17 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.
18 Because this is a spontaneous process: Cell PotentialThe potential voltage a rxn can produce.Reduction potentialsCu e- Cu Eo = .34 VAg+ + e- Ag Eo = .80VSince both rxns are reduction, one must be oxidation, flip it, positive voltage must result from spontaneous rxnsBecause this is a spontaneous process:(Ag+ + e- Ag) x Eo = .80VCu Cu e Eo = -.34 VCu + 2Ag+ Cu Ag Eo = .46 V
19 Because this is nonspontaneous process: Cell PotentialThe potential voltage a rxn can produce.Na+ + e- Na Eo = VNonspontaneous, must end in negative voltage. Flip one to become oxidation.** Fuel Cell!Cl2 + 2e- 2Cl- Eo = 1.36 VBecause this is nonspontaneous process:(Na+ + e- Na) x 2 Eo = V2Cl- Cl e- Eo = V2Na Cl- 2Na + Cl2 Eo = V
21 Rxns that require an energy source to react ElectrochemistryWhen electric voltage is used to produce a redox reaction, it is called electrolysisElectrolytic CellRxns that require an energy source to react
22 BatteriesCar Battery- rechargeable b/c the alternator reverses the ½ rxns and regenerates the reactants.Discharge Cycle Rxn:Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O
23 ElectroplatingAn electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surfaceTypically, an inactive metal is able to be ionized and then deposited on the surface of a more active metal to prevent corrosion.AnodeSilver ions are reduced at the cathode:Ag+ + 1e- → AgSilver atoms are oxidized at the anode:Ag → Ag + + 1e-Cathode
24 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 energyElectrolytic = non-spontaneouselectrical energy → chemical energy
25 ElectrolysisAnode: 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.