1 / 18

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.

marilynjones
Download Presentation

Electrochemistry

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related