Oxidation-reduction (redox) reactions are another important type ofreaction that you will see questions about on the
SAT II Chemistrytest. The test writers will expect you to be able to identify elementsthat are oxidized and reduced, know their oxidation numbers, identifyhalf-cells, and balance redox reactions. The following is a briefoverview of the basics.
Oxidation-Reduction
Oxidation-reduction reactions involve thetransfer of electrons between substances. They take placesimultaneously, which makes sense because if one substance loseselectrons, another must gain them. Many of the reactions we’veencountered thus far fall into this category. For example, allsingle-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. Sinceelectrons are negative, this will appear as an increase in the charge(e.g., Zn loses two electrons; its charge goes from 0 to +2). Metalsare oxidized.
Oxidizing agent (OA): The species that is reduced and thus causes oxidation.
Reduction: The gain of electrons. When anelement 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 andgoes 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 ifthe oxidation numbers of the elements in the reaction change in thecourse of the reaction. We can determine which elements undergo achange in oxidation state by keeping track of the oxidation numbers asthe reaction pro
gresses. You can use the following rules to assignoxidation 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).
- The oxidation state of a monatomic ion is the same as its charge.
- In compounds, fluorine is always assigned an oxidation state of -1.
- Oxygen is usually assigned an oxidation state of -2 in itscovalent 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). - Hydrogen is assigned an oxidation state of +1. Metal hydridesare an exception: in metal hydrides, H has an oxidation state of -1.
- The sum of the oxidation states must be zero for an electrically neutral compound.
- For a polyatomic ion, the sum of the oxidation states must equal the charge of the ion.
Now try applying these rules to a problem.
Example
Assign oxidation numbers to each element in the following:
- H2S
- MgF2
Explanation
- The sum of the oxidation numbers in this compound must be zerosince 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 onH. The sulfur S must have a charge of -2 since there is only one atomof sulfur, and +2 - 2 = 0, which equals no charge.
- F is assigned an oxidation state of -1 (according to rule3), 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 iselectrically neutral.
- This time the net charge is equal to -3 (the charge of thepolyatomic ion—according to rule 7). Oxygen is assigned a -2 oxidationstate (rule 4). Multiply the oxidation number by its subscript: -2
4 = -8. Since there is only 1 phosphorus, just use those algebra skills: P + -8 = -3. Phosphorus must have a +5 charge.
Example
When powdered zinc metal is mixed with iodinecrystals and a drop of water is added, the resulting reaction producesa great deal of energy. The mixture bursts into flames, and a purplesmoke 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.
Explanation
- Assign oxidation numbers to each species. Zn and I2 areboth assigned values of 0 (rule 1). For zinc iodide, I has an oxidationnumber of -1 (group 7A—most common charge), which means that for zinc,the oxidation number is +2.
- Evaluate the changes that are taking place. Zn goes from 0to +2 (electrons are lost and Zn is oxidized). The half-reaction wouldlook like this:
Zn0
Zn2+ + 2e-
And I2 goes from 0 to -1 (it gains electrons and so is reduced). This half-reaction would look like this:
- Here, zinc metal is the reducing agent—it causes the reduction totake place by donating electrons—while iodine solid is the oxidizingagent; 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 occursthrough an external pathway instead of directly between the twoelements. The figure below shows a typical voltaic cell (this onecontains the redox reaction between zinc and copper):
Akx1Or 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 befilled with agar, which contains a neutral salt, or be replaced with aporous cup. Remember that electron flow always occurs from anode tocathode, through the wire that connects the two half-cells, and avoltmeter is used to measure the cell potential in volts.
Batteries are cells that are connected inseries; the potentials add to give a total voltage. One common exampleis 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 willdepend on the half-cells that are involved. Each half-cell has a knownpotential, called its
standard reduction potential (Eº
).The cell potential is a measure of the difference between the twoelectrode potentials, and the potential at each electrode is calculatedas the potential for
reduction at the electrode. That’s why they’re standard reduction potentials, not standard oxidation potentials. Here is the chart:
Y#B6R3AsS On this reduction potential chart, the elements that have the
most positivereduction potentials are easily reduced and would be good oxidizingagents (in general, the nonmetals), while the elements that have theleast positive reduction potentials are easily oxidized and would begood reducing agents (in general, metals). Let’s try a quick problem.
Example
Which of the following elements would be most easily oxidized: Ca, Cu, Fe, Li, or Au?
Explanation
Use the reduction potential chart: nonmetalsare at the top and are most easily reduced. Metals are at the bottomand are most easily oxidized. Lithium is at the bottom of thechart—it’s the most easily oxidized of all. So the order, from mosteasily oxidized to least easily oxidized, is Au, Fe, Cu, Ca, Li.
Example
Which one of the following would be the best oxidizing agent: Ba, Na, Cl, F, or Br?
Explanation
Using the reduction potential chart and thefact that oxidizing agents are the elements that are most easilyreduced, we determine fluorine is the best oxidizing agent.
Electrolytic Cells
While voltaic cells harness the energy fromredox reactions, electrolytic cells can be used to drive nonspontaneousredox reactions, which are also called
electrolysis reactions.Electrolytic cells are used to produce pure forms of an element; forexample, they’re used to separate ores, in electroplating metals (suchas applying gold to a less expensive metal), and to charge batteries(such as car batteries). These types of cells rely on a battery or anyDC source—in other words, whereas the voltaic cell
is a battery, the electrolytic cell
needsa battery. Also unlike voltaic cells, which are made up of twocontainers, electrolytic cells have just one container. However, likein voltaic cells, in electrolytic cells electrons still flow from theanode to the cathode. An electrolytic cell is shown below.
