CH 23

Electrochemistry

23.1 Electrochemical cells

Types of electrochemical cells

Galvanic or VoltaicGalvanic or Voltaic

The ‘spontaneous’ reaction.

Produces electrical energy.ElectrolyticElectrolyticNon-spontaneous reaction.Requires electrical energy to occur.

For reversible cells, the galvanic reaction can occur spontaneously and then be reversed electrolytically - rechargeable batteries.

Voltaic cells

There are two general ways to conduct an oxidation-reduction reaction

Mixing oxidant and reductant togetherMixing oxidant and reductant together

Cu2+ + Zn(s) Cu(s) + Zn2+

This approach does not allow forcontrol of the reaction.

Voltaic cells

Electrochemical cellsElectrochemical cells Each half reaction is put in a separate ‘half cell.’

They can then be connected electrically.

This permits better control over the system.

Spontaneous Reactions

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

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

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

Voltaic cells

Cu2+ + Zn(s) Cu(s) + Zn2+

Zn Cu

Cu2+Zn2+

e- e-

Electrons aretransferred fromone half-cell tothe other usingan external metalconductor.

Electrons aretransferred fromone half-cell tothe other usingan external metalconductor.

Voltaic cells

e- e-

To complete thecircuit, a saltbridge is used

To complete thecircuit, a saltbridge is used

salt bridge

Voltaic cells

Salt bridgeSalt 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.

KClCl- K+

Cl- is releasedto Zn side as Zn is converted to Zn2+

K+ is releasedas Cu2+ isconverted to Cu

Voltaic cells

Zn Zn2+ + 2e-

For our example, we havezinc ion being produced.

This is an oxidation so: The electrode is - the anode - is positive (+).

“AN OX”

Voltaic cells

Cu2+ + 2e- Cu

For our other half cell, we havecopper metal being produced.

This is a reduction so: The electrode is

- the cathode - is negative (-) “RED CAT”

Voltaic Cell

Cell diagrams

Rather than drawing an entire cell, a type of shorthand can be used.

For our copper - zinc cell, it would be:

Zn | Zn2+ (1M) || Cu2+ (1M) |

The anode is always on the left.

| = boundaries between phases

|| = salt bridge

Other conditions like concentration are listed just after each species.

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/MnO2 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

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) + PbO2(s) + 2H2SO4(aq)-----2PbSO4(s) + 2 H2O(l)

Now you write the half reactions that occur at each electrode!!

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 PbSO4 which can fall to the bottom of the battery

Fuel Cell

Idea here is to have a renewable electrode so electrodes don’t wear out

A fuel is used for the oxidationSimplest is the Hydrogen-oxygen fuel

cell- Oxygen is fiels for cathode (reduction), and hydrogen is fuel for anode (oxidation)

Overall reaction: 2H2(g) + O2(g)—2H2O(l)

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.