In an oxidation-reduction reaction, or redox reaction, one or more electrons are transferred from one species to another. For example, in the reaction of sodium metal with chlorine gas to produce sodium chloride salt, electrons are transferred from the sodium metal to the chlorine atoms to create ions, which are held together in the crystal lattice by electrostatic forces. The relatively simple act of transferring electrons can completely change the physical and chemical properties of the substance. Oxidation is the loss of electrons; reduction is gaining electrons. You can remember this using OIL RIG: Oxidation Is Loss, Reduction Is Gain. In the reaction of solid sodium and chlorine gas to produce sodium chloride, sodium is losing electrons, so it is oxidized, and chlorine is gaining electrons, so it is reduced. With the formation of an ionic compound, it's easy to identify which element is gaining electrons and which is losing electrons. With molecular compounds, however, it isn't as intuitive. So we can use oxidation numbers (or oxidation states) to keep track of the transfer of electrons. The oxidation number of an atom is the charge the atom would have if all lone pairs of electrons are considered only on that atom and all electrons in a bond are considered to be on the more electronegative atom. Oxidation numbers help us keep track of electrons, particularly in molecular compounds. We have five rules for assigning oxidation states. These rules are hierarchical, which means if two rules conflict, you apply the rule that's higher up on the list. Our most important rule is that the oxidation state of an atom in a free element is zero. Examples of free elements would be ozone, which is O3, solid iron, or liquid bromine. A free element could be a molecule, but it only contains one element. Our second most important rule is that the oxidation state of a monatomic ion is equal to the charge on the ion. For example, iron 3+ has an oxidation state of +3; nitride, which is N3-, has an oxidation state of -3. Our third most important rule is that the sum of the oxidation states of all atoms in a neutral molecule or formula unit is zero, or polyatomic ion is equal to the charge of the ion. We'll use this rule extensively in our examples. The fourth most important rule is that in their compounds, group 1A metals always have an oxidation state of +1, and group 2A metals always have an oxidation state of +2. Our last rule applies to nonmetals. We assign oxidation states to nonmetals according to this table, which is also hierarchical. First, we assign fluorine an oxidation state of -1; then hydrogen is assigned an oxidation state of +1; then oxygen gets -2; then group 7A, the rest of our halogens, get -1; then group 6A, the rest of oxygen's column, gets -2; and then group 5A gets -3. Let's practice assigning oxidation states. We'll start with methane, CH4. Methane is not a free element or monatomic ion, so the most important rule that applies is our third rule: the sum of the oxidation states of all the atoms in methane is equal to zero, since methane is a neutral molecule. So carbon plus four hydrogen equals zero. From our table of nonmetals, hydrogen has a +1 oxidation state. Solving for carbon, we get -4. So carbon has an oxidation state of -4 and hydrogen has an oxidation state of +1 in methane. Now let's find the oxidation states for carbon and oxygen in carbon dioxide. Our most important rule is that the sum of the oxidation states in a neutral molecule is equal to zero, so carbon plus two oxygen equals zero. From our table of nonmetals, oxygen has an oxidation state of -2. Solving for carbon, we get +4. So how is it that carbon can have a -4 oxidation state in methane and a +4 oxidation state in carbon dioxide? Let's go back to our definition. The oxidation state is the charge an atom would have if all of the lone pairs of electrons are attributed to the atom they reside on, and all of the electrons in the bond are assigned to the more electronegative atom. Let's look at our Lewis structures. In methane, carbon has a single bond with four different hydrogen atoms. Since carbon is more electronegative than hydrogen, carbon is assigned all the electrons from the bonds. A neutral carbon atom has four valence electrons, but this carbon atom is assigned eight valence electrons, so it has a charge of negative four according to the oxidation state. The hydrogen atoms in methane, however, are not assigned any electrons. Since hydrogen normally has one valence electron, the oxidation state of hydrogen is +1. In carbon dioxide, carbon is bonded to oxygen. Since oxygen is more electronegative, all of the electrons in the double bond are attributed to oxygen. Carbon normally has four valence electrons, but in the Lewis structure no electrons are attributed to it in calculating its oxidation state. So it has a +4 oxidation state. Each oxygen, on the other hand, has eight valence electrons attributed to it, so it has an oxidation state of -2. By applying our rules for assigning oxidation states, we can quickly calculate the oxidation states without needing to draw the Lewis structures. Let's try a few more examples. With Cl2, we have two rules that might apply. Chlorine is in our table of nonmetals under group 7A. According to the table, chlorine should receive a -1 oxidation state. However, Cl2 is a free element, so the most important rule that applies is our first rule: the oxidation state of an atom in a free element is zero. With HCl, the most important rule is that the sum of the oxidation states for all atoms in a neutral molecule is zero. So the oxidation state of hydrogen plus chlorine equals zero. Both hydrogen and chlorine are in our table of nonmetals, but hydrogen is higher up, so we assign hydrogen +1 and calculate the oxidation state of chlorine. In HCl, the oxidation state of chlorine happens to be -1, the same value as in the table. But let's determine the oxidation states of chlorine and oxygen in perchlorate. The most important rule that applies is the sum of the oxidation states of all atoms in the polyatomic ion is equal to the charge on the ion. So the oxidation state of chlorine plus four oxygen equals -1. Oxygen and chlorine are both in our table of nonmetals, but oxygen is higher on the list. So oxygen gets its oxidation state of -2, and we calculate the oxidation state of chlorine. Chlorine has an oxidation state of +7. We have to be careful to apply our rules and table hierarchically.