Transcript Oxidation and Reduction - Champlain Valley Union High School

Oxidation and
Reduction
Chapters 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 → Al(aq)+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 → 2Al(aq)+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
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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 → Al(aq)+3
Activity Series of Metals (p. 668)
When the the reaction happens,
electrons move from Al to Cu
• 2Al(s)0 → 2Al(aq)+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
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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•
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Cu+2(aq) + 2e- → Cu(s)
Half Cell- Zn (anode)
Pushes e- to Cu
(cathode)
oxidation
reduction
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 (Eox0) plus
reduction potential (Ered0)
• 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!