electrochemistry n.
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brian-barker

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Electrochemistry
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  1. Electrochemistry The study of the interchange of chemical and electrical energy.

  2. Electrochemical Reactions • All electrochemical reactions involve the transfer of electrons and are therefore, oxidation-reduction reaction. • Electrons are transferred from the reducing agent to the oxidizing agent. • Oxidation is a loss of electrons (increase in oxidation number)-”OIL” • Reduction is a gain of electrons (decrease in oxidation number)-”RIG”

  3. Types of Electrochemical Cells • Galvanic or Voltaic Cells-those in which a spontaneous chemical reaction produces an electrical current that can be used to do work. • Electrolytic Cells-those in which electrical energy from an outside source causes a nonspontaneous reaction to occur.

  4. Components of a Galvanic Cell • Cell-the reacting system • Electrodes-surfaces where the electric current exits or enters 1) anode- electrode compartment in which oxidation occurs. “AN OX” 2) cathode-electrode compartment in which reduction occurs. “RED CAT” • Salt Bridge- U-tube filled with an electrolyte or a porous disk in a tube connecting the two solutions. • Wire-path by which the electrons flow from one compartment to the other. • Electrons flow through the wire from the reducing agent to the oxidizing agent (from the anode to the cathode)

  5. Diagram of Galvanic Cell

  6. Cell Potential • Cell potential (Ecell) or electromotive force (emf) is the “pull” or driving force on the electrons. • The unit of electrical potential is the volt (V) which is defined as 1 joule/coulomb. • Cell potential is measured with a voltmeter.

  7. Standard Reduction Potentials • Reactions in galvanic cells are broken down into half-reactions with each being assigned a reduction potential. • All half reactions are assigned reduction potentials using the standard hydrogen electrode as the reference. (see page 796) • The potentials are all given as reduction processes. • If the process must be reversed (oxidation process), the sign for the potential is reversed. • Since reduction potential is an intensive process (doesn’t depend on the how many times the reaction occurs), the value of the reduction potential is not changed when a half-reaction is multiplied by an integer to balance an equation.

  8. Standard Reduction Potentials (continued) • The more positive the Eo value for a half-reaction, the greater tendency for the half-reaction to occur. • The more negative the Eo value for a half-reaction, the greater tendency for the half-reaction to occur in the opposite direction. • If Eocell> 0 (positive), the forward reaction is spontaneous. • If Eocell < 0 (negative), the forward reaction is not spontaneous and would have to be carried out in an electrolytic cell.

  9. Complete the practice problems on page 797. • A. 0.71 V • B. 0.32 V

  10. Zinc-Copper Galvanic Cell

  11. The Zinc-Copper Cell • Example: Zn(s) + Cu2+(aq)  Zn2+ (aq)+ Cu(s) • Anode: Zn  Zn2+ + 2 e- • Cathode: Cu2+ + 2e-  Cu • Eocell = .337 + .763 = 1.10 V • Line Notation: Zn | Zn2+ | | Cu2+ | Cu (anode is written on the left side and the vertical line represents a phase difference or boundary)

  12. Write the line notation for a galvanic cell consisting of copper (II) and chromium (III) • Cr3+ + 3e- Cr Eo = -0.74 V • Cu2+ + 2e-  Cu Eo = 0.337 V • Copper reduction occurs at the cathode • Chromium oxidation occurs at the anode. • Line Notation: Cr | Cr3+ | | Cu2+ | Cu

  13. Cell Potential, Work, and Free Energy • The work that can be accomplished when electrons are transferred through a wire depends on the “push” behind the electrons. • Potential difference (V) = work (J)/charge (C) • E = -w/q • Work is viewed from the point of view of the system. (Work flowing out of a system is indicated by a minus sign). • In any real, spontaneous process some energy is wasted due to frictional heating-the actual work realized is always less than the calculated maximum.

  14. Electrical Charge • The charge on one mole of electrons is a constant called the faraday (F), which has the value 96,485 coulombs of charge per mole of electrons. • q = nF (n is the number of moles of e-) • w (∆G) = -nFEmax • Solve example 17-3 on page 802.

  15. The Nernst Equation • The Nernst Equation is used to calculate electrode potential and cell potentials for concentrations and partial pressures other than standard-state values. • E = Eo – (2.303 RT/nF ) log Q • E = potential under nonstandard conditions • Eo = standard potential • R= 8.314 • T = temp in Kelvin • n= number of moles of electrons transferred • F = 96,485 C/mole of e- • Q=reaction quotient

  16. Electrolytic Cells • In an electrolytic cell, an outside source of voltage is used to force a nonspontaneousredox reaction to take place. • Oxidation takes place at the anode and reduction takes place at the cathode just as it does in a galvanic cell. • The cell potential in an electrolytic cell < 0. • Electrolytic cells are used in electroplating.

  17. Electrolysis Problems • I = q/t • I = current (amperes, A) • 1 amp = 1C/sec • q = charge (coulombs, C) • t = time (sec) • Once the charge is known, solve the problem as a stoichiometry problem.

  18. Practice Problem #1 • How long must a current of 5.00 A be applied to a solution of Ag+ toproduce 10.5 g of silver metal?

  19. Practice Problem #2 • What mass of Co can be produced from aqueous Co2+ in 1 hour with a current of 15 A?

  20. Practice Problem #3 • A zinc-copper battery is constructed as follows at 25oC: • Zn Zn2+ (0.1M) Cu2+ (2.50M) Cu The mass of each electrode is 200g. • Calculate the cell potential when this battery is first connected. • Calculate the cell potential after 10.0A of current has flowed for 10.0 h. • Calculate the mass of each electrode after 10.0 h • How long can this battery deliver a current of 10.0A before it goes dead?