Oxidation and ReductionChapters 20 & 21

Oxidation vs Reduction

• Oxidation= A substance loses electrons

• Reduction A substance gains electrons

• 2Al(s)0 + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s)

0

• Al(s)0 → AlAl(aq)(aq)

+3+3 aluminum oxno increases

• Cu(aq)+2 → Cu(s)

0 copper oxno decreases

What’s really happening…

• 2Al(s)0 + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s)

0

• 2Al(s)0 → 2AlAl(aq)(aq)

+3+3 + 6e- Al oxno increases

• 3Cu(aq)+2 + 6e- → Cu(s)

0 Cu oxno decreases

• These are called half reactions. Notice the number of electrons lost is the same as the number gained.

What????

• Oxidation= increase in oxygen atoms, increase in oxno, loss in electrons

• Reduction= loss of oxygen atoms, decrease in oxno, gain in electrons

• Remember:• “L.E.O. says G.E.R.” • Loss of electrons oxidation• Gain electrons reduction

Agents

• The oxidizing agent causes oxidation of another substance. Example: copper

• Cu(aq)+2 → Cu(s)

0

• The reducing agent causes reduction of another substance. Example: aluminum

• Al(s)0 → AlAl(aq)(aq)

+3+3

Activity Series of Metals (p. 668)

When the the reaction happens, electrons move from Al to Cu

• 2Al(s)0 → 2AlAl(aq)(aq)

+3+3 + 6e-

• 3Cu(aq)+2 + 6e- → Cu(s)

0

• This electron flow can be measured as voltage!• We will see how later.

Types of Redox Reactions

• Direct Combination:• S + O2 → SO2

• Decomposition:• HgO → 2Hg + O2

• Single Replacement:• Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)

• Cu(s) + 2Ag+(aq) → Cu+

(aq) + 2Ag(s) (net ionic)• But: Cu(s) + ZnCl2(aq) → NR• Cu(s) + Zn+2

(aq) → No reaction (due to relative reactivity rank)

Balancing Redox Equations

• Some equations are are difficult to balance by inspection or trial and error that worked up until now.

• The fundamental principle is that the number of electrons lost in the oxidation process must equal the number of electrons gained in the reduction process.

Electrochemical cells

• Use redox reactions to either produce or use electricity.

Voltaic Cells

• In late 1700’s Italian physician Luigi Galvani twitched frog legs by connecting two metals. Italian scientist Alessandro Volta concluded the two metals in the presence of water produce electricity.

Voltaic Cells

• Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

• Zn(s) → Zn+2(aq) + 2e- oxidation

• Cu+2(aq) + 2e- → Cu(s) reduction

• Half Cell- Zn (anode)

• Pushes e- to Cu

• (cathode)

Voltaic Cell

• Electrons move spontaneously from the anode (-) to the cathode (+)

• The salt bridge allows

• Electrons to move freely

• Without mixing solutions.

Cell Potential

• Ability to move e- through a wire from one electrode to another is the electrical or cell potential. It is measured in volts (v)

• For Example: A Zn-Cu cell with 1 M solutions produces 1.10 volts.

• Here is how it works:

Standard Reduction Potentials (p. 693)

Standard Electrode Potentials

• Ecell = Eoxidation + Ereduction

• Ecell0= sum of the oxidation potential (Eox

0) plus reduction potential (Ered

0)

• The standard state conditions are noted with the 0.

• E0 are determined by measuring half cell potential differences.

• Zn(s) → Zn+2(aq) + 2e- E0

ox = + 0.76 V

• Zn+2(aq) + 2e- → Zn(s) E0 red = - 0.76 V

Calculating Cell Potentials

• Zn(s) → Zn+2(aq) + 2e- E0

ox = + 0.76 V

• Cu+2(aq) 2e- → Cu(s) E0

red = + 0.34 V

• Total Voltage (Ecell) = + 1.10 Volts

• Practice Problems #1 and 2 on P. 696.

That’s it for this

• Electrifying lecture!