This document discusses different types of electrochemical cells including galvanic/voltaic cells which produce electrical energy from spontaneous reactions and electrolytic cells which require electrical energy for non-spontaneous reactions. It describes how voltaic cells work by separating the oxidation and reduction half reactions into two half cells connected by a salt bridge or porous disk to allow ion flow. Common examples discussed include copper-zinc voltaic cells, dry cells like flashlights, lead storage batteries in cars, and hydrogen-oxygen fuel cells.

1. CH 23 Electrochemistry 23.1 Electrochemical cells

2. Types of electrochemical cells Galvanic or Voltaic The ‘spontaneous’ reaction. Produces electrical energy. Electrolytic Non-spontaneous reaction. Requires electrical energy to occur. For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.

3. Voltaic cells There are two general ways to conduct an oxidation-reduction reaction Mixing oxidant and reductant together Cu 2+ + Zn (s) Cu (s) + Zn 2+ This approach does not allow for control of the reaction.

4. Voltaic cells Electrochemical cells Each half reaction is put in a separate ‘half cell.’ They can then be connected electrically. This permits better control over the system.

5. Spontaneous Reactions Will occur if the anode metal is above the cathode metal in the Activity Series chart (pg 678)

6. Voltaic Cell Allessandro Volta (1745-1827) invented the first electrochemical cell- this type was called the voltaic cell. It is a spontaneous reaction- he layered Cu and Zn plates, separated by cardboard: Cu plate had reduction occur, Zn plate had oxidation occur

7. Voltaic Cell A half-cell is one part of the voltaic cell where either oxidation or reduction occurs. A half cell consists of a metal strip immersed in a solution of it’s ions

8. Voltaic cells Cu 2+ + Zn (s) Cu (s) + Zn 2+ Zn Cu Cu 2+ Zn 2+ e - e - Electrons are transferred from one half-cell to the other using an external metal conductor.

9. Voltaic cells e - e - To complete the circuit, a salt bridge is used salt bridge

10. Voltaic cells Salt bridge Allows ion migration in solution but prevents extensive mixing of electrolytes. It can be a simple porous disk or a gel saturated with a non-interfering, strong electrolyte like KCl . KCl Cl - K + Cl - is released to Zn side as Zn is converted to Zn 2+ K + is released as Cu 2+ is converted to Cu

11. Voltaic cells For our example, we have zinc ion being produced. This is an oxidation so: The electrode is - the anode - is positive (+). “ AN OX” Zn Zn 2+ + 2e -

12. Voltaic cells For our other half cell, we have copper metal being produced. This is a reduction so: The electrode is - the cathode - is negative (-) “ RED CAT” Cu 2+ + 2e - Cu

13. Voltaic Cell

14. Cell diagrams Rather than drawing an entire cell, a type of shorthand can be used. For our copper - zinc cell, it would be: Zn | Zn 2+ (1M) || Cu 2+ (1M) | The anode is always on the left. | = boundaries between phases || = salt bridge Other conditions like concentration are listed just after each species.

15. Dry Cell Voltaic cell where the electrolyte is a paste- not a solution Example: flashlight battery ( pg 681) Not a true battery Outer Zn case is anode (oxidation) Carbon (graphite core) rod in center is cathode- but actually reduction occurs w/ MnO 2 found in paste Salt bridge is not needed because of paste prevent cell contents from mixing Alkaline batteries use KOH in paste and this makes it last longer and keeps voltage up

16. Lead Storage Battery A battery is a group of cells connected together A car battery is 6 cells producing 2V each for a total of 12 V The cathode is lead(IV) oxide and the anode is Pb. Dilute sulfuric acid is the electrolyte Overall reaction is: Pb (s) + PbO 2(s) + 2H 2 SO 4(aq) -----2PbSO 4(s) + 2 H 2 O (l) Now you write the half reactions that occur at each electrode!!

17. Lead Battery Car battery’s are recharged when the car runs- the reaction occurs in reverse- but this reverse reaction is nonspontaneous and so the car’s generator supplies the energy to drive the reaction. Eventually the battery dies- electrodes lose so much PbSO 4 which can fall to the bottom of the battery

18. Fuel Cell Idea here is to have a renewable electrode so electrodes don’t wear out A fuel is used for the oxidation Simplest is the Hydrogen-oxygen fuel cell- Oxygen is fiels for cathode (reduction), and hydrogen is fuel for anode (oxidation) Overall reaction: 2H 2(g) + O 2(g) —2H 2 O (l)

19. Fuel Cell You write the anode and cathode half-cell reactions. Advantage: cheap fuel, only “pollutant”- water which is drinkable Used in spacecraft and some military applications- some cars; expensive and takes room.