Transcript Lecture 9. Redox chemistry

Lecture 9. Chemistry of
Oxidation-Reduction Processes
Prepared by PhD Halina Falfushynska
Oxidation-Reduction Reactions
• Often called “redox” reactions
• Electrons are transferred between the
reactants
– One substance is oxidized, loses electrons
• Reducing agent
– Another substance is reduced, gains
electrons
• Oxidizing agent
• Oxidation numbers change during the reaction
LEO says GER
Lose Electrons = Oxidation
1
0
Na  Na  e

Sodium is oxidized
Gain Electrons = Reduction
0

1
Cl  e  Cl
Chlorine is reduced
• Rules for assigning oxidation numbers
1. Elements (uncombined) are 0.
Al, N2, He, Zn, Ag, Br2, O2, O3
2. Oxidation numbers must sum to the overall
charge of the species.
SO42 = 2 (O is usually 2 so….)
? + 4(2) = 2
Solve: ?  8 = 2
? = + 6 (S)
Guidelines for Assigning Oxidation Numbers
is 1 and for KO2 is ½.
Assign oxidation numbers for all elements
in each species
MgBr2
Mg +2, Br 1
ClO2
Cl +1 , O 2
Oxidation Numbers on the Periodic Table
(most common in red)
Copyright McGraw-Hill 2009
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• Displacement reactions
– A common reaction: active metal
replaces (displaces) a metal ion from a
solution
Mg(s) + CuCl2(aq)  Cu(s) + MgCl2(aq)
– The activity series of metals is useful in
order to predict the outcome of the
reaction.
• Balancing redox reactions
– Electrons (charge) must be balanced as
well as number and types of atoms
– Consider this net ionic reaction:
Al(s) + Ni2+(aq)  Al3+(aq) + Ni(s)
– The reaction appears balanced as far as
number and type of atoms are
concerned, but look closely at the charge
on each side.
Al(s) + Ni2+(aq)  Al3+(aq) + Ni(s)
– Divide reaction into two half-reactions
Al(s)  Al3+(aq) + 3e
Ni2+(aq) + 2e  Ni(s)
– Multiply by a common factor to
equalize electrons (the number of
electrons lost must equal number of
electrons gained)
2 [Al(s)  Al3+(aq) + 3e ]
3 [Ni2+(aq) + 2e  Ni(s) ]
– Cancel electrons and write balanced net ionic
reaction
2Al(s)  2Al3+(aq) + 6e
3Ni2+(aq) + 6e  3Ni(s)
2Al(s) + 3Ni2+(aq)  2Al3+(aq) + 3Ni(s)
Predict whether each of the following will
occur. For the reactions that do occur,
write a balanced net ionic reaction for
each.
- Copper metal is placed into a solution of silver
nitrate
Cu (s)
+ 2 Ag (aq)
Cu 2+ (aq) +
2 Ag(s)
- A gold ring is accidentally dropped into a solution
of hydrochloric acid
No reaction occurs, gold is below
hydrogen on the activity series.
•
Combination Reactions
– Many combination reactions may also
be classified as redox reactions
– Consider:
Hydrogen gas reacts with oxygen
gas
2H2(g) + O2(g)  2H2O(l)
Identify the substance oxidized and
the substance reduced.
• Decomposition reactions
– Many decomposition reactions may also
be classified as redox reactions
– Consider:
Potassium chlorate is strongly heated
2KClO3(s)  2KCl(s) + 3O2(g)
Identify substances oxidized and
reduced.
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• Disproportionation reactions
– One element undergoes both oxidation
and reduction
– Consider:
16
• Combustion reactions
– Common example, hydrocarbon fuel
reacts with oxygen to produce carbon
dioxide and water
– Consider:
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Reaction of Cu and Zn2+ ions
Gets Smaller ->
<- Gets Larger
Cell Notation
Zn (s) + Cu2+ (aq)
Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
Zn (s)| Zn+2 (aq, 1M)|K(NO3) (satur)|Cu+2(aq, 1M)|Cu(s)
Salt bridge
anode
cathode
Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)
K(NO3)
Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt
Electrochemical Cells
The difference in
electrical potential
between the anode and
cathode is called:
•cell voltage
• electromotive force (emf)
• cell potential
0
Cell
E
E
0
oxidation
E
0
reduction
Standard Electrode Potentials
Standard reduction potential (E0) is the voltage
associated with a reduction reaction at an electrode
when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction
2e  2H (1 M)
E0  0 V
Standard hydrogen electrode (SHE)
H2 (1 atm)
Determining if Redox Reaction is Spontaneous
• + E°CELL ; spontaneous
reaction
• E°CELL = 0; equilibrium
• - E°CELL;
nonspontaneous
reaction
More positive E°CELL ;
stronger oxidizing agent or
more likely to be reduced
Relating E0Cell to DG0
ECell
work

ch arg e
Units
work, Joule
charge, Coulomb
Ecell; Volts
Faraday, F; charge on 1 mole eF = 96485 C/mole
DG = -nFEcell
Relating EoCELL to the
Equilibrium Constant, K
DG0 = -RT ln K
DG0 = -nFE0cell
-RT ln K = -nFE0cell
J 

8
.
31

(298K )
RT 
molK 

 0.0257
C
F
96485
mole
0
Cell
RT

ln K
nF
0
Cell
0.0257
0.0592

ln K 
log K
n
n
E
E
Corrosion – Deterioration of Metals by
Electrochemical Process
Corrosion – Deterioration of Metals by
Electrochemical Process
Corrosion – Deterioration of Metals by
Electrochemical Process
Corrosion
•Damage done to metal is costly to
prevent and repair
•Iron, a common construction metal often
used in forming steel alloys, corrodes by
being oxidized to ions of iron by oxygen.
•This corrosion is even faster in the
presence of salts and acids, because
these materials make electrically
conductive solutions that make
electron transfer easy
Corrosion
•Luckily, not all metals corrode easily
•Gold and platinum are called noble
metals because they are resistant to
losing their electrons by corrosion
•Other metals may lose their electrons
easily, but are protected from corrosion by
the oxide coating on their surface, such as
aluminum – Figure 20.7, page 636
•Iron has an oxide coating, but it is not
tightly packed, so water and air can
penetrate it easily
Corrosion
•Serious problems can result if bridges,
storage tanks, or hulls of ships corrode
•Can be prevented by a coating of oil,
paint, plastic, or another metal
•If this surface is scratched or worn away,
the protection is lost
•Other methods of prevention involve the
“sacrifice” of one metal to save the second
•Magnesium, chromium, or even zinc
(called galvanized) coatings can be applied