Reduction-Oxidation Reactions
Any reaction where electrons are transferred between reactants is called an reduction-oxidation
reaction or redox reaction. These are reactions where one substance wants an electron so badly
that it takes it away from another substance. Whether or not it succeeds depends on who it meets.
For example, consider the reaction
FeCl2(aq) + CeCl4(aq) → FeCl3(aq) + CeCl3(aq)
What happened here? If we remove the spectator ions and write the net ionic equation we find:
Fe2+(aq) + Ce4+(aq) → Fe3+(aq) + Ce3+(aq)
Ce4+ took an electron from Fe2+! This is an oxidation/reduction reaction. In this example, Fe2+ is
oxidized and Ce4+ is reduced. The charge of Fe went from +2 to +3, that is, it lost an electron.
This process is called oxidation.
Oxidation: The loss of an electron by a substance.
Likewise, the charge of Ce went from +4 to +3, that is, it gained an electron. This process is
called reduction.
Reduction: The gain of an electron by a substance.
Half-Reactions
Oxidation/reduction reactions are important because we can exploit them as a way of generating
electrical current. For example, we know that Ce4+ will pull an electron away from Fe2+ when we
mix the two in solution. The trick to making a battery is to find a way to make Ce4+ pull an
electron from Fe2+ when they are not mixed together in a single solution.
In order to do this we set up two 1/2 reactions in separate beakers and connect them with a salt
bridge. The salt bridge electrically connects the two beakers, but prevents Fe2+ and Ce4+ from
mixing.
BY DR. J.J. GONGDEN: www.vitalchemist.com
You need two 1/2 reactions to make a reaction, so, we add them together.
Fe2+ → Fe3+ + eCe4+ + e- → Ce3+
__________________________________
Fe2+ + Ce4+ →Fe3+ + Ce3+
Notice that electrons on both sides of the half-reactions must cancel each other out when added
together.
What about other atoms and molecules. How do you know if one chemical substance is strong
enough to take an electron from another? We simply refer to a list known as the activity series.
Oxidation States
Oxidation numbers provides a means of keeping track of electrons in redox reactions. For some
elements, the oxidation number is just another way of stating what its most stable cation or anion
will be. In other cases, it is not so obvious, so we have rules.
Rules for assigning oxidation numbers
1. The oxidation number for elements is always zero. For example, Na(s), O2(g), C(s) all
have zero oxidation numbers.
2. The oxidation number of monoatomic ions is the same as their charge. You already
know this one. This means that for Na+, the oxidation number is +1 and for Cl-, the
oxidation number is -1.
3. Oxygen is assigned a -2 oxidation number in covalent compounds. This refers to
compounds such as CO, CO2, SO2, and SO3. There is an exception to this rule, and it is in
peroxides, such as H2O2. Here, each O in the O22- group has a -1 oxidation number.
4. Hydrogen is assigned a +1 oxidation number in covalent compounds. This refers to
compounds such as HCl, NH3, and H2O.
5. In binary compounds, the element with the greatest attraction to electrons gets the
negative oxidation number. In other words, the most electronegative of the pair gets the
negative number. For example, in HF, F is more electronegative and thus has a -1
oxidation number. In NH3, the N atom is more electronegative and has a -3 oxidation
number.
6. The sum of the oxidation numbers is zero for a neutral compound and equal to the
ion's charge for an ionic species. For example, in H2O, a neutral species, H is +1 and O
is -2, and the sum of the two is 0. For CO32-, each O is -2 and C is +4, and the sum is -2.
BY DR. J.J. GONGDEN: www.vitalchemist.com
Assign oxidation numbers to the atoms in SF6.
Since this is a binary compound let's first start with rule 5. We know that F has a greater
attraction to electrons than S does, therefore we give it the negative oxidation number, which in
this case will be -1 for F. To assign sulfur its oxidation number, we go to rule 6. This is a neutral
compound, so the sum of the total oxidation numbers must be zero. Since there are 6 F atoms,
each with a -1 oxidation state then the sulfur must have a +6 oxidation number to balance out the
fluorine atoms.
BY DR. J.J. GONGDEN: www.vitalchemist.com