Electrochemistry

Voltaic Cell (or Galvanic Cell) The energy released in a spontaneous redox reaction can be used to perform electrical work. A voltaic cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants. In a voltaic cell, chemical energy is changed to electrical energy.

Anode/Cathode The two solid metals that are connected by the external circuit are called electrodes. The electrode at which oxidation occurs is called the anode. The electrode at which reduction occurs is called the cathode.

Half Cells The voltaic cell is thought of as being comprised of two "half-cells." One cell is where oxidation occurs and the other is reduction.

Half Reactions Anode (oxidation half-reaction): Zn (s)→Zn +2 (aq) + 2e- Cathode (reduction half-reaction): Cu +2 (aq) +2e- →Cu (s) Determining half reactions will be shown later.

OIL – RIG Remember the acronym OIL – RIG Oxidation Is Loss of electrons Reduction Is Gain of electrons

Red-Ox Reaction In a red-ox reaction, one substance must be oxidized and another substance must be reduced. The substance that is oxidized is the reducing agent. The substance that is reduced is the oxidizing agent.

Cell Operation Electrons become available when the zinc metal is oxidized at the anode. The electrons flow through the external circuit to the cathode, where they are consumed as Cu +2 is reduced. Because zinc is oxidized in the cell, the zinc electrode loses mass, and the concentration of the Zn +2 solution increases as the cell operates. As the same time, the Cu electrode gains mass, and the Cu +2 solution becomes less concentrated as the Cu +2 is reduced to Cu (s).

Cell Operation As the voltaic cell operates, oxidation of Zn introduces additional Zn +2 ions into the anode compartment. Unless a means is provided to neutralize this positive charge, no further oxidation can take place. At the same time, the reduction of Cu +2 at the cathode leaves an excess of negative charge in solution in that compartment.

Salt Bridge Electrical neutrality of the system is maintained by a migration of ions through the porous glass disc (salt bridge) that separates the two compartments. A salt bridge consists of a U-shaped tube that contains an electrolyte solution whose ions will not react with other ions in the cell or with the electrode materials.

Salt Bridge As oxidation and reduction proceed at the electrodes, ions from the salt bridge migrate to neutralize charge in the cell compartments.

Ion Migration Anions migrate toward the anode and cations toward the cathode. No measurable electron flow will occur through the external circuit unless a means is provided for ions to migrate through the solution from one electrode compartment to another, completing the circuit.

Electron Flow In any voltaic cell the electrons flow from the anode through the external circuit to the cathode. Because the negatively charged electrons flow from the anode to the cathode, the anode in a voltaic cell is labeled with a negative sign and the cathode with a positive sign.

Oxidation Numbers Remind yourself how to determine oxidation numbers. Oxidation numbers show what the charge of each atom would be (in a molecule or ion), if each atom were an ion. Go to the text and revisit the Rules for Assigning Oxidation Numbers section.

Balancing Red-Ox Equations Obey the Law of Conservation of Mass Gain and loss of electrons must be balanced

Half Reactions Cu + 2Ag + → 2 Ag + Cu +2 Determine the oxidation number for each substance in the reaction Write two half reactions based on one substance being oxidized and one being reduced.

Half Reactions Cu → Cu e-oxidation 2e- + 2Ag + → 2 Agreduction Half reactions show the number of electrons gained or lost by the substance, standard redox equations do not. Go to ChemReview packet to practice half reaction writing.

Review Sample Ex: 20.4 Try Practice Exercises

Cell EMF What is the driving force that pushes the electrons through an external circuit in a voltaic cell? An oxidizing agent in one compartment pulls electrons through a wire from a reducing agent in the other compartment. The pull (or driving force) on the electrons is called the cell potential ( E cell ) or electromotive force (emf) of the cell. Or, electrons flow from the anode to the cathode (in a voltaic cell) because of a difference in potential energy.

Potential Energy of Electrons The potential energy of electrons is higher in the anode than in the cathode. Therefore electrons spontaneously flow through an external circuit from the anode to the cathode. For any cell reaction that proceeds spontaneously (i.e. voltaic cell), the cell potential will be positive.

Standard Conditions Standard Conditions: * 1M concentrations for reactants and products in solution * 1 atm pressure (for gases) * 25 ˚C Under Standard Conditions the emf is called the standard emf or the standard cell potential (E ˚ cell ) Keep in mind that the superscript ˚ denotes standard- state conditions.

Cell Potential The cell potential of a voltaic cell depends on the particular cathode and anode half-cells involved. Standard potentials have been assigned to each individual half-cell, and then use the half-cell potentials to determine E ˚ cell. The cell potential is the difference between two electrode potentials (anode and cathode). The potential associated with each electrode is chosen to be the potential for reduction to occur at that electrode.

Standard Reduction Potentials Standard electrode potentials are tabulated for reduction reactions, so they are called standard reduction potentials ( E ˚ red ). E ˚ cell = E ˚ red (cathode) - E ˚red (anode) Or E ˚ cell = E ˚ red (cathode) + E ˚ox (anode) The hydrogen half-reaction was used as a reference for all of the other half-reactions. The standard reduction potential for the hydrogen half-reaction is assigned a standard reduction potential of exactly 0 V. It is also called the standard hydrogen electrode. (SHE)

Calculating EMF We will use standard reduction potentials to calculate the emf of a voltaic cell. Because electrical potential measures potential energy per electrical charge, standard reduction potentials are intensive properties. Changing the stoichiometric coefficient in a half-reaction does not affect the value of the standard reduction potential.

Calculating EMF Try sample exercise 20.5

E ˚ red and spontaneity The more positive the E ˚ red value for a half- reaction, the greater the tendency for the reactant of the half-reaction to be reduced, and, therefore, to oxidize another species. Try sample exercise 20.8

EMF and Free Energy Change ΔG = - nFE n is a positive # that represents the # of e- transferred in the reaction F is Faraday’s constant (this constant is the quantity of electrical charge on 1 mole of electrons) 1F = 96,500 C/mol = 96,500 J / V mole The units for ΔG are J/mole

EMF and Free Energy Change ΔG = - nFE If n and F are positive values, a positive value of E leads to a negative ΔG. Remember that a negative ΔG indicates a spontaneous reaction. The equation can be altered slightly if the reactants and products are in their standard states: ΔG˚ = - nFE˚

Concentration and Cell EMF Remember ΔG = ΔG˚ + RT ln Q Combining ΔG = ΔG˚ + RT ln Q and ΔG = - nFE You have the Nernst equation: E = E˚ – RT ln Q nF

Nernst Equation The Nernst equation can be expressed two ways: E = E˚ – RT ln Q nF E = E˚ – 2.303RT log Q nF

Nernst Equation Variation The equation can be simplified if the cell is run under standard conditions: E = E˚ – V log Q n * Temp. is 298 K

Use the Nernst Equation to work through the example on the bottom of page 773 Try the sample exercise and practice exercise

Electrolysis This process is the opposite of a voltaic cell.

Ch 20 Problems 6, 7, 9, 14, 24, 25 a,b, 27, 29, 31, 36, 43, 44, 46, 49, 51, 52, 56, 59