REDOX AND OXIDATION POTENTIALS
In any oxidation-reduction reaction, the same number of electrons must be donated as are accepted. In this lesson, you will
explore some of the consequences of this rule.
The tendency of electrons to be donated is a measure of the strength of a reducing agent. If an element or molecule will
readily donate an electron, it is called a strong reducing agent. It will be oxidized in a redox reaction as it reduces the other
reactant. An element which possesses a strong ability to receive an electron is a strong oxidizing agent. It will be reduced in
a redox reaction. As it oxidizes the other reactant, it receives electrons and is itself reduced. Oxygen is a very strong
oxidizing agent and readily available in our environment. Hence, the name of the process "oxidation."
Examine the data table below. You will notice that all the reactions shown are oxidation reactions because the reactant is
oxidized as it loses one or more electrons. These reactants are considered the reducing agents because they donate the
electron(s). Starting at the top, with the strongest reducing agent, the list provides the relative tendency of the reducing
agents to donate on electron.
At the bottom of the list you will find that fluorine has the least tendency to give up an electron. Fluorine may therefore be
considered to be the weakest reducing agent or the strongest oxidizing agent. That is, in a redox reaction involving F, it is
more likely that this element will receive an electron by oxidizing the other reactant.
How has the relative position of these reactions been determined? This has been accomplished by measuring the voltage
produced in a Galvanic cell similar to that of the figure below showing the electrochemical cell. Each reaction is allowed to
take place under controlled conditions. The electron transfer which takes place occurs through a wire. The resulting voltage
associated with the electric current is called the oxidation potential (Eo). The relative position of Eo is actually more important
than the absolute value. All the values of oxidation potential may simply be considered relative to the reaction for H 2, which
has been assigned an arbitrary value of zero volts.
Relevance of Eo. Lithium has the greatest oxidation potential on this table. This helps us to understand that because Li is
the strongest reducing agent, it yields the greatest current (flow of electrons) when allowed to react in an electrochemical
Galvanic cell. The voltage measured is directly proportional to the current produced. The oxidation reaction shown for Li will
therefore have the highest potential of occurring. This table helps us to predict which reactants will be oxidized and which
will be reduced when they are brought together. Remember the zinc and copper chemical equation?
Zn0 + Cu+2
Zn+2 + Cu0
From the data table, it can be seen that Zn is a stronger reducing agent then Cu (higher on the list). Alternately, you may
determine that Cu is a stronger oxidizing agent than Zn. This means that when both are allowed to react with one another,
zinc may be expected to be oxidized by the stronger oxidizing agent of Cu. Copper may be expected to be reduced by the
stronger reducing agent of Zn. Likewise Zn would be expected to provide the oxidation reaction because its oxidation
potential (Eo) is the greater of the two.
Summary of the data table:
Any Reducing Agent will reduce an agent below it.
Any Reactant on the list will act as an Oxidizing Agent for any agent above it.
Take a close look at what happens in an electrochemical cell (see the preceding figure). The term electrochemical implies
that we are converting chemical energy into electrical energy. Look again at the Zn - Cu+2 system. At the zinc electrode,
electrons are released and oxidation occurs. The electrode where oxidation occurs is called the anode. Electrons are gained
at the copper electrode and reduction occurs. The electrode where reduction occurs is called the cathode. As electrons
leave the cell from the anode, positively charged Zn+2 ions are produced. Electrical neutrality is maintained. Because
negative ions always drift toward the anode, they are called anions. Since positive ions drift away from the anode toward the
cathode, they are called cations. The Zn+2 ion drifts away into solution, but electrons remain in the zinc electrode. They
move up the zinc electrode, through the wire, and enter the copper cathode. At the surface of the rod, the electrons
encounter Cu+2 ions in solution. The electrons react with the Cu+2 to give neutral copper atoms, which remain on the rod as
copper metal. The reaction continues until either the zinc bar is depleted or the Cu +2 ions in solution are all used up. Writing
separate "half-reaction" equations for the two separate processes is often convenient:
(1) Zn0
Zn+2 + 2 e-(oxidation)
(2) Cu+2 + 2 e-
Cu0 (reduction)
The net reaction (3) is found by adding the two half-reactants.
Notice that the electrons on each side of the equation cancel. If they do not cancel in the half-reactions, as in the Cu0 - Ag+
reaction, (below) we must first balance the electrons and then add the half-reactions. That is:
Later in the section you will learn how the balancing of electrons will enable you to write balanced equations for more difficult
reactions.
Oxidation numbers. We are now ready to turn our attention away from oxidation potentials and take a look at the
bookkeeping of electrochemical reactions. Oxidation numbers are arbitrary numbers assigned to an atom within a molecule.
These numbers are useful in balancing complex equations. Just as rules are needed in bookkeeping for procedures to
follow, a chemist needs a few rules to assist him in his accounting of electron charge. The following rules will assist you in
assigning oxidation numbers to various atoms.
1. The oxidation number of a monatomic ion is equal to the charge on the ion. For example: Na +1 (Group 1), Mg+2 (Group 2),
and so on.
2. The oxidation number of any substance in its atomic state is zero. For example, Cu 0, H20, Cl20, Pb0.
3. The oxidation number of hydrogen in a compound is +1.
4. The oxidation number of oxygen in a compound is -2 (except in O2, O3 gases, and in peroxides).
5.The halogens (F, Cl, Br, I) are assigned an oxidation number of -1.
6. The Laws of Conservation of Charge and Atoms must be followed.
7. Other oxidation numbers are assigned so that the sum of charges is zero.
Example 1: Assign oxidation numbers to the atoms in K3PO4.