CH 23 Electrochemistry 23.1 Electrochemical cells
Types of electrochemical cells
- The ‘spontaneous’ reaction.
- Produces electrical energy.
- Non-spontaneous reaction.
- Requires electrical energy to occur.
- For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.
- 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.
- Each half reaction is put in a separate ‘half cell.’ They can then be connected electrically.
- This permits better control over the system.
- Will occur if the anode metal is above the cathode metal in the Activity Series chart (pg 678)
- 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
- 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
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.
Voltaic cells e - e - To complete the circuit, a salt bridge is used salt bridge
- Salt bridge Allows ion migration in solution but prevents extensive mixing of electrolytes.
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
- It can be a simple porous disk or a gel saturated with a non-interfering, strong electrolyte like KCl .
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 -
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
- 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
- Other conditions like concentration are listed just after each species.
- Voltaic cell where the electrolyte is a paste- not a solution
- Example: flashlight battery ( pg 681)
- Outer Zn case is anode (oxidation)
- Carbon (graphite core) rod in center is cathode- but actually reduction occurs w/
- 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
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
- 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!!
- 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
- 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)
- 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.