Oxidation Numbers: An Ecological Perspective

Assigning oxidation numbers for the purposes of predicting behavior is one of 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 might have to live less greedily and use technology more wisely in converting energy in the absence of combustion.

Science Isn’t There to Control Nature

Although I enjoy most of Jacob Bronowski’s attempts at linking science to literature and to the fine arts, can you guess what reservations I have regarding his introduction in “The Creative Process” ? (Unfortunately I could not find an online version; I have a paper-reprint of his essay which appeared in a September 1958 issue of Scientific American.)

“The most remarkable discovery made by scientists is science itself. The discovery must be compared in importance with the invention of cave-painting and of writing. Like these earlier human creations, science is an attempt to control our surroundings by entering into them and understanding them from inside. And like them, science has surely made a critical step in human development which cannot be reversed. We cannot conceive a future society without science.”

Once we have created models that help us understand an aspect of nature, there’s nothing about the scientific approach that obliges us to control our surroundings. It’s with technology that we make such an attempt. Yes, we cannot conceive of a society without science for a couple of reasons, but neither has anything to do with manipulating nature. (1) It is highly compatible with our curious nature. Its concepts, although neither flawless nor immutable, give us a better approximation of the truth than what we would obtain from commonsense notions, wild guesses and superstition.

(2) Some of the problems created by a combination of human nature, social structures, science and technology need science as part of the solution.

But simply because we seem to gain control over nature, both out of necessity and greed, is no reason to include that characteristic in a definition of science.

Ironically, while successfully erasing the false dichotomy between the arts and sciences, Bronowski falls into the trap of imagining a new polarity. He imagines that a civilization which expresses itself mostly in contemplation values no creative activity. At the other extreme, we have Western scientists, scholars and artists whose outlook is supposedly always active.

As Bronowski points out, through metaphors artists and scientists do create unity in what is diverse. To use examples from a field that I know, in each of the 8 million known species of organisms, there are always carbon-based compounds present. Each carbon atom in those life forms can form several different molecular orbitals with other carbons, hydrogen, oxygen and other elements.  But thermodynamically the most stable arrangement consistently involves four bonds for each carbon. Now granted, it took a mentally and experimentally active approach on the part of many people to create those concepts, but now that we have them , what compels us to quickly synthesize far more compounds on a massive scale without knowing their full impact on our bodies and on the rest of nature? It’s not inherent to science.

Couldn’t we should just slow down and value knowledge more for its own sake? Should we not also express ourselves more frequently in contemplation? Can’t we look at a dead leaf and be content with knowing that when it continues to decompose in the spring, it will allow nitrates, phosphates and water to be retained by carbon-based soil a little longer and let new life recycle them? And instead of rushing to paint them or to look for more compounds in the stages of an oak leaf’s decomposition, why not extend the enjoyment of those transient moments of early morning light interacting with its frost-coated surface? In fact, the most creative people do pause more often. They do it not only for the sake of incubation but for the sake of existential pleasures.

Nature has set up a museum in the park this morning. It charges no admission and needs no curator. Ironically, after contemplating them, I could not resist the temptation to photograph them!

The Joys of Walking Across An Icy Field

Education has met its goals, not necessarily when it has landed you a dream job—which I think is an illusion for the vast majority of people who have slaved in the past and who are working now— but when it can intensify the sensual and intellectual pleasures of the simplest acts of life—like walking to work.

Same field as being described but from another day.

Weather-wise, we have had an erratic month of February in Montreal, with more than the usual cycles of freezing and melting. One day, after a morning of snow and an afternoon of freezing rain, the snowscape was glazed with a thick, milky ice, thick enough to support the weight of a toddler. It was cold for a few days afterwards, but subsequent rain transformed the veneer. I was reminded that with every subzero drop, the forms of snow and ice, like the size of all crystals, depend on how quickly the temperature drops and on the impurities and imperfections that seed them. During my 2 km-walk that morning, I experienced different textures and densities.  This is because sources of dust, pools of water,  their depth and amount of surface exposed to air are all variable and not every spot is equally affected by wind and footsteps.

Over two thousand five hundred footsteps that gave me an uplifting return on by body’s investment in adenosine triphosphate, ATP. (For the uninitiated, ATP is not a drug.  It is the currency of cells, the facilitator for all of our energy-requiring reactions. It’s what chloroplasts produce when photosynthesizing before creating glucose, and it is what our cells create when oxygen breaks down the metabolites of that same sugar. )

As I felt the unmonotonous sequence of pressures around my boots, each different area that I walked on created a unique sound  Since childhood, my favourite sound of that type is that of thin ice shattering above an air pocket. I was also reminded that the frictional coefficient of naturally-formed ice varies significantly. That morning no one else braved the -22 oC windchill factor. So with no one watching, I was a free 54-year old, giving myself a little run and testing to see how far I’d glide on various sections of the ice field.

The way light interacted with all the surfaces also accentuated their differences. There were sparkles from icy particles acting as tiny prisms; there were lakes of yellow-orange as the rising sun caught expanses of smooth surfaces; and other parts of the frozen field glistened with different hues. Every hue corresponded to a different frequency, suggesting a unique interaction of matter with light energy.

Another benefit of immersing myself into the walk was that it lifted the weight of the thoughts about the oncoming day. It also dissipated any of the usual worries that the condensation of water in my breath was beginning to accumulate and cool my neck-warmer. If anything, the journey was far too short. I was tempted to turn 180 degrees and repeat the walk with even more attention to detail…Oh but, wait, I said to myself. The fact that I did not indulge in an extended trek was not a wasted opportunity. After all, there was the journey home that afternoon.