Chapter 19
 Oxidation and reduction reactions involve the transfer of
electrons
 These reactions can be identified through the assigning
oxidation states
 The concept of oxidation states was previously discussed in
Ch 7
 They are an informal way of keeping track of the distribution
of electrons in molecules, compounds, or ions
Elements can exist in a number of different oxidation states.
Assign an oxidation state to sulfur in each of the following
species.
 H 2S
 S8
 SO2
 SO42 S2O3
 S2O32-
 Oxidation – reaction in which an atom or ion experiences an
increase in oxidation state
 Oxidation – the loss of electron density
2Na(s) + Cl2(g) → 2NaCl(s)
Na → Na+ + e The oxidation state of sodium has changed from 0, its
elemental state, to the +1 state of its ion.
 Since its oxidation number is increasing, the Na is said to
have been oxidized.
 Reduction – reaction in which an atom or ion experiences a
decrease in oxidation state
 Reduction – the gain of electron density
2Na(s) + Cl2(g) → 2NaCl(s)
Cl2 + 2e- → 2Cl The oxidation state of chlorine has changed from 0, its
elemental state, to the -1 state of its ion.
 Since its oxidation number is decreasing, the Cl2 is said to
have been reduced.
For each reaction, determine whether oxidation or reduction
has occurred and then balance the reaction using electrons.
1. K → K+
2. O2 → O23. Fe2+ → Fe3+
4. Mn7+ → Mn2+
Which of the following represent redox reactions?
1. 2KNO3(s) → 2KNO2(s) + O2(g)
2. H2(g) + CuO(s) → Cu(s) + H2O(l)
3. H2(g) + Cl2(g) → 2HCl(g)
4. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
5. SO3(g) + H2O(l) → H2SO4(aq)
For each reaction
that is a redox
reaction, identify
the element which
is oxidized and
which is reduced.
A reducing agent is a substance that causes
another substance to be reduced.
 Reducing agents themselves must lose
electrons and are oxidized
An oxidizing agent is a substance that
causes another substance to be oxidized.
 Oxidizing agents themselves must gain
electrons and are reduced
The relative strength of reducing and
oxidizing agents can be determined using
Table J on your reference tables.
 Active metals lose electrons easily
 Active metals are easily oxidized
 The more active the metal, the better
reducing agent it is
Li is the strongest reducing agent listed on
Table J
What about Li+?
Li+ has already lost its valence electron and
does not want to lose any more. It could only
gain one back (act as an oxidizing agent) but
it will not be very effective in doing so.
What would be the most effective oxidizing
agent on Table J?
 Active non-metals gain electrons easily
 Active non-metals are easily reduced
 The more active the non-metal, the better
oxidizing agent it is
F2 is the strongest oxidizing agent listed on
Table J
What about F-?
F- has already gain a valence electron and
does not want to gain any more. It could only
lose one (act as a reducing agent) but it will not
be very effective in doing so.
Using Table J, predict whether or not the
following chemical reactions would occur.
Would Al be oxidized by Ni2+?
Would Cu be oxidized by Ag+?
Would Pb be oxidized by Na+?
Would Al be oxidized by Cu2+?
Exhausted after arriving home from a long
day at the lab, you remember something.
You left an iron spatula sitting in a solution of
Sn(NO3)2 that you made in preparation for an
experiment the next morning.
Should you go back to remove it or will it be
fine until the morning?
Fe(s) + Sn(NO3)2(aq) → ?
Fe(s) + Sn2+(aq) → ?
Fe(s) + Sn2+(aq) → Fe2+(aq) + Sn(s)
Electrochemistry is the branch of chemistry that deals with
electricity related applications of oxidation-reduction (redox)
reactions
Fe(s) + Sn2+(aq) → Fe2+(aq) + Sn(s)
This reaction is a spontaneous redox reaction and the electrons
are transferred from one substance (Fe) to the other substance
(Sn2+) directly in the solution.
If, however, we can separate the substances that are competing for
electrons, we can harness the electrons transferred to produce
electricity.
Oxidation
Fe(s) → Fe2+(aq) + 2e-
(-)
Reduction
Sn2+(aq) + 2e- → Sn(s)
Fe(s) + Sn2+(aq) → Fe2+(aq) + Sn(s)
Anode
site of
oxidation
Salt bridge
Maintains
electrical
neutrality
by
permitting
the flow of
ions
(+)
Cathode
site of
reduction
Sn
Fe
Fe2+
Sn2+
A voltaic cell (sometimes called a galvanic cell or battery), makes use of a
spontaneous redox reaction to convert chemical energy to electrical energy
Zinc and Copper
Some redox reactions do not occur spontaneously, but can be
driven by electrical energy.
Electrolysis – the process in which an electric current is used to
produce a redox reaction
 Electrical energy to chemical energy
2Al3+(aq) + 3Cu(s) → 2Al(s) + 3Cu2+(aq)
 This was a redox reaction that we predicted would not be
spontaneous
 Could be driven by application of an electric current
A rechargeable cell combines the
redox chemistry of both voltaic and
electrolytic cells
When a rechargeable battery operates
as a voltaic cell, it converts chemical
energy to electrical energy
When the cell is recharged, it operates
as an electrolytic cell, converting
electrical energy into chemical energy
Standard 12V automobile battery is a
set of six rechargeable cells
Anode reaction:
Pb(s) + SO42-(aq) → PbSO4(s) + 2eCathode reaction:
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- →
PbSO4(s) +2H2O(l)
A car’s battery provides the electrical
energy needed to start the engine
Once the car is running, the half
reactions are reversed by a voltage
produced by the alternator
Batteries can be recharged as long as
all of the components needed are
present and the reaction is reversible
Metals such as copper, silver, and
gold are difficult to oxidize
In an electrolytic cell, these inactive
metals form ions at the anode that
can be easily reduced at the
cathode
This allows solid metal from one
electrode to be deposited on the
other electrode
This process is called electroplating
+
What significance does the label “Voltaic Cell” have?
The term voltaic cell means that chemical energy will be converted into
electrical energy. In other words, electrons will flow in the spontaneous
direction.
When the switch is closed, in which direction will electrons flow?
Pb to Ag or Ag to Pb
What is the purpose of the salt bridge?
To maintain electrical neutrality (complete the circuit) by permitting
the migration of ions.
When the switch is closed, in which direction will negative ions
flow?
Pb(s) → Pb2+(aq) + 2e-
Which electrode will act as the anode?
Write the half-reaction that occurs.
Pb(s) → Pb2+(aq) + 2e-
Ag+(aq) + e- → Ag(s)
Which electrode will act as the cathode?
Write the half-reaction that occurs.
Pb(s) → Pb2+(aq) + 2e-
Ag+(aq) + e- → Ag(s)
Write the overall redox reaction that occurs in this electrochemical cell.
Pb(s) +2Ag+(aq) → 2Ag(s) + Pb2+(aq)
Pb(s) → Pb2+(aq) + 2e-
Ag+(aq) + e- → Ag(s)
What happens to the mass of each electrode as the cell discharges?
The mass of the Pb electrode decreases and the mass of the Ag
electrode increases.
Ag+(aq) + e- → Ag(s)
Eventually the cell potential decreases to zero. Why does this happen?
The redox reaction has reached chemical equilibrium.