C.W. BAKER HIGH SCHOOL
Chemistry CORE VIII – Oxidation-Reduction
Prepared by M. Foster
After reviewing this study
guide you may want to test your knowledge with the following sample
questions and answers.
1.
Oxidation – Reduction reactions –
oxidation-reduction reactions are chemical reactions in which electrons are
transferred from one reactant to the other.
These reactions involve the two processes:
a.
Oxidation –is the process by which a substance loses
electrons. The name comes from the fact
that substances frequently lose electrons when they react with oxygen; however,
oxygen is not required for the oxidation process.
b.
Reduction – is the process by which a
substance gains electrons.
c.
LEO roars GER – many students
keep the definitions of oxidation and reduction straight by using this
“nonsense” saying where LEO stands for “Loss electrons – oxidation and GER
stands for “Gain electrons –reduction.
d.
REDOX Reactions – this term is frequently used in
referring to oxidation reduction reactions because both oxidation and reduction
must take place at the same time in order to have a reaction. Students should also remember that the number of
electrons gained in the reduction process must be equal to the number of electrons
lost in the oxidation process.
2.
Determining whether a chemical reaction is a REDOX reactions – Since redox
reactions involved the transfer of electrons it is possible to judge whether a
reaction is redox or not by looking at the oxidations states of the reactants
and products. If the oxidation state of
any reactant atom changes as it goes to a product, the reaction is a redox
reaction.
a.
Oxidation states – these numbers indicate the number of
electrons which have been “gained” or “lost” by an atom. For example, an atom which has neither lost
or gained any electrons would have a zero oxidation state. Atoms which have lost electrons are assigned
positive oxidations states (a +2 oxidation state means the atom has lost 2 electrons) while atoms
which have gained electrons are assigned negative oxidation states ( a –2 oxidation
state means the atom has gained 2 electrons)
b.
Rules for determining oxidation states – there are
five rules which can be utilized in determining the oxidation states of atoms.
i.
The oxidation state of any element in its standard state (uncombined with
another element) is zero.
ii.
The oxidation state of any monatomic atom is equal to the charge on
the atom.
iii.
The oxidation state of hydrogen in a compound is almost always +1
iv.
The oxidation state of oxygen in a compound is almost always –2.
v.
The sum of the oxidation states for a particle equals the charge on
the particle (for a molecule = 0, for an ion = charge on the ion).
c.
Single replacement reactions – are always redox
reactions because they involve the transfer of electrons.
Ex. 2Na +
H-OH ---à NaOH +
H2
In this case,
the oxidation state of the sodium goes from zero to +1, while the oxidation
state of the hydrogen goes from +1 to zero.
d. Double replacement
reactions – are never redox reactions because there is no change in the
oxidation states of the products and reactants (and hence no transfer of
electrons).
3.
Half reactions – redox reactions can be broken up into
half-reactions. One half reaction
represents oxidation, while the other half reaction represents reduction.
a.
Reduction half reaction – the reduction
half reaction shows the substance which is gaining the electron(s) and the
substance which is formed after the electrons are gained. The electrons always appear as reactants in
the reduction half reaction.
Ex. Cu+2 +
2 e- à Cuo
b.
Oxidation half reaction – the oxidation
half reactions shows the substance that is losing the electrons and the
substance that is formed after the electrons have been lost. The electrons always appear as a product
in the oxidation half reaction.
Ex. Zn à Zn+2 +
2 e-
4.
Balancing redox reactions – redox reactions
can be balanced by writing the half reactions.
After writing the half reactions, the number of electrons lost in the
oxidation half reaction should be compared to the number of electrons gained in
the reduction half reaction. If they
are not equal, find the least common multiple.
Then multiply the entire half reactions by the appropriate factor to
insure that the number of electrons lost equals the number of electrons
gained.
Ex. Given the overall reaction Zn
+ Cr+3 à Cr + Zn+2
a.
Write the reduction and oxidation half reactions.
Reduction: Cr+3 + 3 e-
à Cr
Oxidation: Zn à Zn+2 +
2 e-
b.
Since the number of electrons gained are not equal to the number
of electrons lost (3 does not equal 2), determine the least common multiple
(6). Multiply the half reactions by the
appropriate number to make the number of electrons lost equal to the number of
electrons gained.
2 Cr+3 +
6 e- à 2 Cr
3
Zn à 3
Zn+2 + 6 e-
c.
Add the half reactions together (cancelling the electrons since
they are on opposite sides of the arrow) to get the balanced equation.
3 Zn
+ 2 Cr+3 à 3 Zn+2 + 2 Cr
5.
Oxidizing Agents/Reducing Agents
a.
Oxidizing agents are the species that cause some other substance to lose
electrons (to be oxidized).
b.
Reducing agents are the species that cause some other substance to gain electrons (to
be reduced).
c.
Example: in the reaction H2 + Cl2 à 2 HCL the hydrogen atoms
lose electrons, causing the Chlorine atoms to gain electrons. Since the hydrogen causes the chlorine to
gain electrons (to be reduced), the hydrogen is the reducing agent. In the same reaction, by accepting
electrons from the hydrogen, the chlorine causes the hydrogen to lose electrons
(to be oxidized. That makes the
chlorine the oxidizing agent.
6.
Types of Electrochemical Cells
a.
Voltaic Cell – redox reaction that is utilized to
produce electrical energy. (Ex. Batteries)
i.
A voltaic cell can be compared to a hydro-electric power
plant. In the power plant the energy of
the water is harnessed as the water flows downhill in terms of energy. In a voltaic cell, the energy of the
electrons is harnessed as they flow “downhill” in terms of energy.
ii.
Voltaic cells are spontaneous (just as water flows downhill
spontaneously)
b.
Electrolytic Cell – redox reaction that uses electrical
energy to produce chemical energy. (Ex. Electroplating, re-charging batteries)
i.
To continue the analogy to water, this is like pumping water
uphill, it requires energy. In an
electrolytic cell the electrons are “pumped” uphill by an power source.
ii.
Electrolytic cells are non-spontaneous.
7.
Parts of an electrochemical cell
a.
Salt Bridge – part of a voltaic cells which
prevents positive or negative charge from building up at the electrodes.
b.
Cathode – in either a voltaic or electrolytic cell, the electrode
where reduction occurs (Cathode and reduction both begin with a consonant.)
c.
Anode – in either a voltaic or electrolytic cell, the electrode
where oxidation occurs (Anode and oxidation both begin with a vowel).