Redox and Electrochemistry
Oxidation-reduction (redox) reactions are another important
type of reaction that you will see questions about on the SAT II
Chemistry test. The test writers will expect you to be able to identify
elements that are oxidized and reduced, know their oxidation numbers, identify
half-cells, and balance redox reactions. The following is a brief
overview of the basics.
Oxidation-reduction reactions involve the transfer of
electrons between substances. They take place simultaneously, which
makes sense because if one substance loses electrons, another must
gain them. Many of the reactions we’ve encountered thus far fall
into this category. For example, all single-replacement reactions
are redox reactions. Before we go on, let’s review some important
terms you’ll need to be familiar with.
Electrochemistry: The study of the interchange
of chemical and electrical energy.
Oxidation: The loss of electrons. Since electrons
are negative, this will appear as an increase in the charge (e.g.,
Zn loses two electrons; its charge goes from 0 to +2). Metals are oxidized.
Oxidizing agent (OA): The species that is
reduced and thus causes oxidation.
Reduction: The gain of electrons. When an
element gains electrons, the charge on the element appears to decrease,
so we say it has a reduction of charge (e.g., Cl gains one electron
and goes from an oxidation number of 0 to -1). Nonmetals are reduced.
Reducing agent (RA): The species that is
oxidized and thus causes reduction.
Oxidation number: The assigned charge on
an atom. You’ve been using these numbers to balance formulas.
Half-reaction: An equation that shows either
oxidation or reduction alone.
Rules for Assigning Oxidation States
A reaction is considered a redox reaction if the oxidation
numbers of the elements in the reaction change in the course of
the reaction. We can determine which elements undergo a change in
oxidation state by keeping track of the oxidation numbers as the
reaction progresses. You can use the following rules to assign oxidation
states to the components of oxidation-reduction reactions:
The oxidation state of an element is zero,
including all elemental forms of the elements (e.g., N2,
P4, S8, O3).
oxidation state of a monatomic ion is the same as its charge.
compounds, fluorine is always assigned an oxidation state of -1.
is usually assigned an oxidation state of -2 in its covalent compounds.
Exceptions to this rule include peroxides (compounds containing
the group), where each oxygen is assigned an
oxidation state of -1, as in hydrogen peroxide (H2O2).
is assigned an oxidation state of +1. Metal hydrides are an exception:
in metal hydrides, H has an oxidation state of -1.
sum of the oxidation states must be zero for an electrically neutral compound.
a polyatomic ion, the sum of the oxidation states must equal the
charge of the ion.
Now try applying these rules to a problem.
Assign oxidation numbers to each element in the following:
The sum of the oxidation numbers in this compound
must be zero since the compound has no net charge. H has an oxidation
state of +1, and since there are two H atoms, +1 times 2 atoms =
+2 total charge on H. The sulfur S must have a charge of -2 since
there is only one atom of sulfur, and +2 - 2 = 0, which equals no
is assigned an oxidation state of -1 (according to rule 3), and
there are two atoms of F, so this gives F a total charge of -2.
Mg must have a +2 oxidation state since +2 - 2 = 0 and the compound
is electrically neutral.
time the net charge is equal to -3 (the charge of the polyatomic
ion—according to rule 7). Oxygen is assigned a -2 oxidation state
(rule 4). Multiply the oxidation number by its subscript: -24 = -8. Since there is only 1 phosphorus,
just use those algebra skills: P + -8 = -3. Phosphorus must have
a +5 charge.
When powdered zinc metal is mixed with iodine crystals
and a drop of water is added, the resulting reaction produces a
great deal of energy. The mixture bursts into flames, and a purple
smoke made up of I2 vapor is produced from
the excess iodine. The equation for the reaction is
Zn(s) + I2(s)ZnI2(s) + energy
Identify the elements that are oxidized and reduced, and
determine the oxidizing and reducing agents.
Assign oxidation numbers to each species. Zn and I2 are
both assigned values of 0 (rule 1). For zinc iodide, I has an oxidation
number of -1 (group 7A—most common charge), which means that for
zinc, the oxidation number is +2.
the changes that are taking place. Zn goes from 0 to +2 (electrons
are lost and Zn is oxidized). The half-reaction would look like
Zn0Zn2+ + 2e-
And I2 goes from 0 to -1 (it gains
electrons and so is reduced). This half-reaction would look like
Here, zinc metal is the reducing agent—it causes the
reduction to take place by donating electrons—while iodine solid
is the oxidizing agent; iodine solid accepts electrons.
Voltaic (or Galvanic) Cells
Redox reactions release energy, and this energy can be
used to do work if the reactions take place in a voltaic cell. In
a voltaic cell (sometimes called a galvanic cell),
the transfer of electrons occurs through an external pathway instead
of directly between the two elements. The figure below shows a typical
voltaic cell (this one contains the redox reaction between zinc
As you can see, the anode is the electrode
at which oxidation occurs; you can remember this if you remember
the phrase “an ox”—“oxidation
occurs at the anode.” Reduction takes place
at the cathode, and you can remember this with the
phrase “red cat”—“reduction
occurs at the cathode.” An important component
of the voltaic cell is the salt bridge, which is a
device used to maintain electrical neutrality; it may be filled
with agar, which contains a neutral salt, or be replaced with a
porous cup. Remember that electron flow always occurs from anode
to cathode, through the wire that connects the two half-cells, and
a voltmeter is used to measure the cell potential in volts.
Batteries are cells that are connected in
series; the potentials add to give a total voltage. One common example
is the lead storage battery (car battery), which has a Pb anode,
a PbO2 cathode, and H2SO4 electrolyte
is their salt bridge.
Standard Reduction Potentials
The potential of a voltaic cell as a whole will depend
on the half-cells that are involved. Each half-cell has a known
potential, called its standard reduction potential (Eº).
The cell potential is a measure of the difference between the two
electrode potentials, and the potential at each electrode is calculated
as the potential for reduction at the electrode. That’s
why they’re standard reduction potentials, not standard oxidation
potentials. Here is the chart:
On this reduction potential chart, the elements that have
the most positive reduction potentials are easily
reduced and would be good oxidizing agents (in general, the nonmetals),
while the elements that have the least positive reduction potentials
are easily oxidized and would be good reducing agents (in general,
metals). Let’s try a quick problem.
Which of the following elements would be most easily oxidized:
Ca, Cu, Fe, Li, or Au?
Use the reduction potential chart: nonmetals are at the
top and are most easily reduced. Metals are at the bottom and are
most easily oxidized. Lithium is at the bottom of the chart—it’s
the most easily oxidized of all. So the order, from most easily
oxidized to least easily oxidized, is Au, Fe, Cu, Ca, Li.
Which one of the following would be the best oxidizing
agent: Ba, Na, Cl, F, or Br?
Using the reduction potential chart and the fact that
oxidizing agents are the elements that are most easily reduced,
we determine fluorine is the best oxidizing agent.
While voltaic cells harness the energy from redox reactions,
electrolytic cells can be used to drive nonspontaneous redox reactions,
which are also called electrolysis reactions. Electrolytic
cells are used to produce pure forms of an element; for example,
they’re used to separate ores, in electroplating metals (such as
applying gold to a less expensive metal), and to charge batteries
(such as car batteries). These types of cells rely on a battery
or any DC source—in other words, whereas the voltaic cell is a
battery, the electrolytic cell needs a battery.
Also unlike voltaic cells, which are made up of two containers,
electrolytic cells have just one container. However, like in voltaic
cells, in electrolytic cells electrons still flow from the anode
to the cathode. An electrolytic cell is shown below.