Assigning oxidation numbers for the purposes of predicting behavior is one the key concepts in chemistry. Although there are countless benefits to teaching kids basic literacy and numeracy, if it was all done solely for them to understand oxidation numbers, it would still be worth it!
It initially appears strange to students that an oxidation number is not always grounded in reality. In certain cases, however, they are the same as ionic charges, which are real. And whether real or imaginary, oxidation numbers are extremely useful.
Let’s begin with ions. The nucleus of atoms has a positive charge, which exerts an attractive force for oppositely charged electrons. The nucleus contends with distance between itself and electrons and with screening by other electrons. After all that, some atoms fall into categories of extremes. There are some that are really good at exerting a strong force between its nucleus and its most distant, so called valence electrons, and those that are fairly feeble. Put together a member of each different group, and it doesn’t get much to generate a reaction. Examples include table salt which you can make by gently heating the low-melting sodium metal and placing it in a beaker filled with the yellow-green and poisonous chlorine gas. Or you can flatten soft potassium metal into a disk and use it to cover a few crystals of pure iodine solid. A gentle blow with a hammer will form the other additive of table salt, potassium iodide, which is included in the salt we buy to prevent goiter.
In each of the reactions, the metal atom loses an electron to the non-metal, either chlorine or iodine. The metal in the newly formed salt goes from being neutral to +1, the charge resulting from having lost a negative electron. Because it mimics what substances do in oxygen’s presence, we say that the sodium has been oxidized from 0 to +1. Another way of viewing this is that neutral atoms have an equal number of protons(+) and electrons(-). But after losing one electron, the sodium atom, for example, now has 11 protons and only 10 electrons. The sum of those charges, 11 + (-10) = +1. That is the new oxidation number of sodium in NaCl or of potassium in the salt potassium iodide. For the chloride or iodide ions, the opposite is true. The excess of electrons gives each of their atoms a charge of -1. We say that the chlorine has been reduced, in this case form 0 to -1.
In our story so far, we have covered two rules of assigning oxidation numbers. An atom in its neutral or free (not bonded to another kind of atom) state is assigned an oxidation number of zero. Metal ions and non metal ions, the members of the extreme groups, have oxidation numbers equal to their respective charges.
Why does it matter? When most metals go from a free state to a charged one the process of bonding to a nonmetal takes them to a more favorable thermodynamic state. What does that mean? In reacting, atoms strive for both disorder and for the bottom of the hill in terms of potential energy, which is associated with the attractive force we described. They can’t always get both, but if at least one of the favorable conditions offsets an unfavorable one, it will still make the reaction spontaneous. In this case, making a crystalline salt doesn’t create disorder. But the act of combining an intermediate gaseous metal ion with a gaseous non metal releases an awful lot of heat, causing the overall reaction to be a downhill process.
To make this less abstract, let us use a practical example: the rust that forms on railroad tracks, supporting beams or cars. The main compound involved is Fe2O3. The oxidation number of oxygen in compounds is usually -2. Since the entire compound, Fe2O3 is neutral, solving this simple equation 2Fe +3(-2) = 0 reveals iron’s oxidation state to be +3. Both the +2 and +3 states of iron in compounds are more stable than that of the zero state found in the free iron of steel. To move from zero to +3, all that free iron has to do is lose electrons to an atom who has the opposite problem. That atom is oxygen.
Free oxygen only exists on earth because plants use the energy of sunlight to produce sugars. In so doing they use pigment-centers that lose electrons when absorbing light, but neutrality is restored when water molecules split up to return electrons.In so doing, H2O molecules not only create a proton gradient that’s used to invest the energy of sunlight, but they also produce free oxygen. But each atom in the oxygen molecule can accommodate two more electrons in its valence shell—hence oxygen’s oxidation number of -2 —hence the fact that oxygen created either by stars or photosynthesis ends up as either water or in the main compounds of the earth’s crust: silicates, iron oxide and aluminum oxide.
When we create iron from ore in the planet’s crust we chase out the oxygen with heat and coal and return electrons to the the iron ions. To prevent it from losing electrons again afterwards—and to delay it from reaching its thermodynamic destiny— we have to either paint or wax its surface in an almost vain attempt to block out oxygen. A more durable alternative is to mix it with other metals whose oxidized coating seriously slows down oxygen’s intrusion. Such metals include nickel, chromium, and zinc.
If we let too much iron rust, we get caught up in having to produce more of the free element. The starting materials, rust and ore, will not run out soon, but the heat and carbon required for its reduction depend on fossil fuels, which when consumed, yield carbon dioxide. If we consume too much stainless steel, their production also has repercussions since nickel, chromium and zinc all have to be reduced as well. Their ores are all oxides or sulfide equivalents.
Now we come to assigning oxidation numbers to covalent compounds, where electrons are shared; they are not lost and taken as they are in ionic compounds such as salt and rust. The -2 oxidation number is assigned to the oxygen in CO2, even though it doesn’t represent a true charge. Consistent with that is an oxidation number of +4 for carbon. It turns out that the creation of CO2 through combustion, respiration or decomposition puts the carbon in a molecule that’s in a gaseous and higher state of entropy than any solid or liquid form containing carbon. Moreover, the formation of a pair of tight C=O bonds found in the molecule also releases a fair amount of energy. Figuratively speaking, producing CO2 is a free ride. But CO2 is a greenhouse gas whose concentration jumps dramatically when we oxidize fossil fuels on a massive scale!
Over millions of years, in the absence of the electron thief, oxygen, it took lots of energy in the form of geological pressure to form compounds like CH4(methane), C8H18(octane), and C-aggregates(coal). The oxidation number of carbon in those compounds is -4 , -9/4(an average), and 0, respectively, all lower than the +4 in carbon dioxide. To raise the oxidation state, it takes the same strategy that works for neutral iron: yield electrons to oxygen, even though this time the atoms engaged share the electrons. The end-result is that when burnt, the molecules “gladly” return heat and CO2 in proportion to how many carbon atoms are in each molecule, with methane having the least and coal being cursed with the most.
I was demonstrating to my students that the Mn in potassium permanganate (KMnO4) is in a highly oxidized state, +7 to be exact. This makes it a proficient electron thief. In the presence of glycerin, C3H8O3, beholder of a carbon in a low oxidation state relative to CO2, a reaction usually takes place within a minute that the two are in contact. But this time the reaction was so slow to activate that it actually came to fruition in another class, one hour later.
The reason for this strange occurrence might have boiled down to the fact that the KMnO4 powder I used was far finer than normal. In principle coarse powder reacts more slowly. But my reaction was far slower than what I usually observed from powder that had less surface area per volume. If a powder is too fine, its higher surface area might imply that the additional exposed atoms could obtain their electrons beforehand by snatching them from moisture in the air inside the container. (Something similar occurs if you try reacting old powdered zinc with acid. The reaction is slower than what’s obtained with zinc nuggets!) This feat of having a oxidation-reduction reaction with H2O, rare among man-made chemicals, is possible because KMnO4 is a stronger oxidizing agent than oxygen. Oxygen is created from the oxidation of water in a reaction similar to what occurs in photosynthesis.
To verify this theoretical hypothesis, my wife suggested that I grind the already fine powder with a mortar and pestle just prior to adding the glycerin. This would get some of the moisture-reduced coating off the powder and expose more pure KMnO4. Sure enough, within seconds the same powder and glycerin that had taken so long to react now erupted and produced a beautiful violet-colored flame. (The color is created by the presence of potassium ion.)
In the same way that I was motivated to reduce manganese, we need more members of society to work together with algae and terrestrial plants. Inadvertently, in our absence they succeed in balancing the amount of carbon dioxide in the atmosphere. They do so by reducing the oxidation state of carbon , converting CO2 into sugars and fatty, nucleic and amino acids, which make other forms of life possible. They work against volcanic eruptions, lightning-induced fires, respiration and decomposition, all of which serve to place more of the oxidized form of carbon into the atmosphere. What photosynthesizers cannot do, because there aren’t enough of them, is offset our overzealous oxidation of carbon. We have to live less greedily and use technology more wisely in converting energy in the absence of combustion.