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Electrochemistry

Learn about the setup of a voltaic cell and how to calculate cell potential in electrochemical reactions. Practice problems and battery types are also covered.

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Electrochemistry

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  1. Electrochemistry • Electrochemical Cell – an apparatus that uses redox reactions to produce electrical energy. • Voltaic Cell – a type of electrochemical cell that converts chemical energy to electrical energy with a spontaneous redox reaction.

  2. Set Up • We will use the reaction: Zn(s) + Cu2+ Zn2+ + Cu(s) • We will place 1M ZnSO4 and a strip of zinc on one beaker. • We will place 1 M CuSO4 and a strip of copper in the other beaker.

  3. Set Up • In order for electrons to be exchanged, a wire is connected to each metal piece. • Since zinc in the strip is being oxidized, there will be a build up of positive ions (Zn+2) in solution as electrons leave. • The zinc strip loses mass over time.

  4. Set Up • In the copper beaker, since electrons are entering, there will be a build up of negative ions in the solution. (Cu+2 in solution becomes Cu) • The copper strip gains mass over time.

  5. Set Up • A salt bridge is used to allow ions to flow from one side to the other. • It contains a soluble salt that is contained by a “plug” such as agar gel. • Ions can move through plug, but solutions do not mix. • Anions enter the cell where oxidation occurs. • Cations enter the cell where reduction occurs.

  6. Voltaic Cell • The electrode where oxidation takes place is the anode. • The electrode where reduction takes place is the cathode. • Which metal becomes oxidized or reduced is determined by it’s reduction potential. • Reduction potential is the tendency of an substance to gain electrons. This is measured in the units of volts. • The more negative the number the more likely it will be oxidized, the more positive the more likely it will be reduced.

  7. Reduction Potentials

  8. Calculating Cell Potential • All reduction potentials are measured against hydrogen (which can be oxidized or reduced by the substance.) • To measure the cell we created we first look at the copper.

  9. Calculating Cell Potential • Now to measure what is happening at the zinc electrode. • Since copper has the more positive value, it will be reduced and the zinc will be oxidized.

  10. Calculating Cell Potential • To calculate the overall potential: E0cell= E0red– E0ox • Therefore for our cell: 0.342V-(-0.762V) = +1.104 V • A positive number indicates it occurs spontaneously

  11. Calculate the cell potential of the following: A. Chromium in Cr+3 solution and copper in Cu+2 solution. B. Tin in Sn+2 solution and iodine in I- solution. A. E° = +1.086 V B. E° = +0.6730 V Practice Problems

  12. Battery • Battery- can contain a single cell or packages of several cells • Small batteries (household) contain a single cell. • Large batteries (car) contains many cells that can conduct more current. These are lead-acid batteries.

  13. Types of Batteries • Primary batteries – produce electricity by using a redox reaction that is not reversible. • Once reactants are gone, battery is thrown out. • Typical alkaline batteries oxidize powdered zinc. • Secondary batteries – are rechargeable because redox reactions are reversible. • These batteries are usually made of nickel and cadmium (NiCad).

  14. Types of Batteries • Lithium batteries – use lithium because it is lightest in mass of all metals and has lowest standard reduction potential. • This allows the battery to last much longer.

  15. Some Vocabulary • Electroplating: using electrical current to cause ions in solution to form a metal layer on a surface. • One Ampere: (A) amount of electrical current equal to 1 coulomb of charge per sec. • One Coulomb: (C) amount of charge that passes a point when 1 amp of current flows for 1 sec. (1A = 1C/sec) • One Faraday: (F) amount of electricity that passes per 1 mole of electrons (1F=96485C or 1 mole e- =96485C)

  16. Let’s try some math • A current of 2.50 amps is passed through a solution of Ni(NO3)2 for 2.00 hours. What mass of nickel is deposited at the cathode? • Ans: 5.48 g

  17. Another Problem • If you wish to convert 1.00 grams of Au+3 ions to solid Au, how long must you electrolyze the solution if the current passing through the circuit is 2.00 amps? • Ans: 735 sec

  18. Let’s do one more! • After 9.50 minutes at 5.50 amps, all of the Fe+3 ions were plated out of 600.0 mL of solution. What was the original concentration of Fe+3 in solution? • Ans: 0.0181 M

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