Human minds are a funny thing. Take mine for instance. When I first heard of special relativity in a college physics course, I found it refreshing, but unfortunately not powerful enough to encourage me to pursue it beyond our subject’s superficial introduction. Years later, I went back to the derivation of the gamma factor and that of E = mc2, and I found it as inspiring as the Pink Floyd’s Meddles album that played in my head while I sat in physics class.
Similarly, a friend of mine who developed mathematical models to test new designs of airplanes wondered why I had aborted my math education while majoring in chemistry. In my old age I have gone back to learning math, almost on a daily basis (albeit at a snail’s pace) while that same retired friend is more interested in pursuing his passion for music, probably because math reminds him of work.
Given that I will probably bore him again with another math story; given that my wife will refuse to listen to it outright, while my daughter would only take a polite interest in it while killing the minutes during her walk to the lab— I will tell you about it instead.
Consider the following infinite sum, S, of the form
In the denominator consecutive odd numbers are raised to the same power,d . As we alternately add and subtract more and more terms, the series converges to a fixed number. For example, for the power of 1, ( d = 1):
the first four terms give a result of 0.8349….. If you use a total of 8 terms, we get 0.81309…. Eighty terms produces 0.78848… When we add and subtract 800 terms, the first two decimals don’t change as the series yields 0.78571… The third decimal becomes steady as we move to 1500 decimals. With 15000 terms, we get 0.78541… which looks like π/4 = 0.78540. If we use different odd powers for the sums, an interesting pattern emerges:
Notice that the power of π is always equal to d. Less obvious is that the denominator equals 2d+1 times (d-1)! For example, for d = 1, the denominator equals (21+1) (1-1)! = 4(1) = 4. For d = 5, it equals (25+1) (5-1)! = 64(4*3*2*1) = 1536, and so on.
There is a more succinct way of expressing the sum by using the symbol sigma:
For any whole number d value, odd or even, there is a Dirichlet beta function that’s defined as
So according to our table,
To express our pattern for odd powers of d, then we define x = 2n+ 1 , which means that n= (d-1)/2.
For those familiar with calculus there’s another beautiful relationship between the Dirichlet beta function and an integral. It’s :
From a statistical point of view, oxygen is the sports superstar of the elements. It is the only element to appear in the top 3 by weight percentage. It’s the third most abundant element in the universe; second on earth and life’s first. There are life forms that can get by without diatomic oxygen, but all life depends on the organic compounds and water that can only be formed with oxygen atoms. Here we will begin with the gas that no land animal or plant can do without.
I have opened my share of dry cells and alkaline batteries to reveal their inner workings to myself, and then eventually to students. After sawing off the batteries’ top, there is a characteristic smell. It has reminded me of my first childhood attempt at making the essential element oxygen from my parents’ hydrogen peroxide (H2O2) and an ingredient common to both zinc and the old ammonium-based batteries: manganese dioxide. This fine black powder catalyzes the chemical breakdown of each pair of hydrogen peroxide molecules into a pair of water molecules and one diatomic molecule of oxygen.
You can check for oxygen’s odorless presence with a glowing wood splint. If the concentration in a test tube is high enough, you not only see the splint characteristically burst into a flame, you even get a nice popping sound. In simpler reactions, regardless of whether oxygen interacts with metals or non-metals, the compounds have less potential energy than the reactants. Since energy is conserved, the excess energy is liberated in the form of light and/or heat. But for living aerobic cells, some of the energy from oxidized food molecules is invested in making other compounds which eventually facilitate non-spontaneous reactions. Instead of wreaking havoc like a forest fire after a lightning storm, the subsequent anabolic reactions decrease the internal entropy of the organism, sustaining it. We often hear that oxygen “burns” food. In comparing cellular respiration to combustion, there are indeed common starting and end- products, but the former has other intermediates—a whole other story in between, a difference between life and destruction.
Why do biological food molecules exist? Life can use energy sources to build larger molecules. Later, it evolved photosynthesis, a way to use light as its energy source. In doing so, electrons are emitted from its pigment-centers, solar cells far more sophisticated than the ones we have invented. These electrons help bond smaller molecules into glucose. But that would not in itself be energetically feasible unless the photosynthetic organism could tap into the bulk of the solar energy, which is not contained in the electrons.
Oxygen enters the picture. About 3 billion years ago, in prokaryotes, photosynthesis not only involved electrons leaving pigments, but in order to make pigments available for excitation again, electrons had to return to them. Each of the four electrons, delivered one at a time, came from a pair of water molecules that dissociated to also produce 4 positive ions of hydrogen and a molecule of oxygen. The oxygen slowly filled the atmosphere, which paved the way for the evolution of animals. The charged particles concentrated on one side of a membrane, storing potential energy and allowing the bulk of the energy to be used for the so-called dark reactions of photosynthesis.
Figure 1 The oxygen-evolving-complex. (a) Ca2+ , oxygen and various states of oxidation of Mn make up the inorganic cluster at the center of the proteins. (b) A model for how the oxidation state of manganese ions change. Each boxed S represents a different combination of Mn charges. https://www.sciencedirect.com/science/article/abs/pii/S1367593115000022[/caption
The oxygen-evolving-complex (OEC) ( Figure 1) is an enzyme with an imbedded cluster of Mn4O5Ca. Notice how manganese, the atom and battery-ingredient involved in the catalytic breakdown of hydrogen peroxide makes another appearance. Sitting in one of the pigment-systems of plants and algae, OEC is the precise location of where water molecules break up, deliver their electrons to chlorophyll and yield acid and oxygen gas. But how the enzyme exactly operates is still a big biochemical mystery. It’s certainly more sophisticated than the way we electrolyze water into hydrogen gas and oxygen.
i. Oxygen and fire
It’s hard to discuss oxygen without remembering that along with fuel and heat, it is one of the three parts of the fire triangle. Remove any of the trio and the fire is extinguished. Let’s explore the chemistry of fire and oxygen’s role by examining campfires.
Figure 2. Fire is an interaction of energy and matter, including free radicals and diatomic carbon. On the right, birch bark. Pictures by author.
Birch bark, which is found in many parts of the U.S., Canada, Europe and China, is a great way to start a campfire. Rich in terpenoids, the paper-thin material ignites easily. The heat it releases provides enough activation energy to set small twigs ablaze, which of course should be placed in a tee-pee arrangement to let in more oxygen. All of this should take place in a pit surrounded by stones, not to let wind take heat away from the young fire and not to burn the forest down.
The hues of a flame are rarely constant for a second, a hint that something complex is occurring within them. There is a set of chain mechanisms involving intermediate molecules that are needed for subsequent steps. Many of the in-between products are radicals, reactive molecules with unpaired electrons. Radicals are often created in the high temperature regions of the flame, but they diffuse back into the colder regions where they are needed to generate the final products. To reveal more details, laser-based investigative techniques have been used so as not to disturb the flame. Even when burning as simple a molecule as diatomic hydrogen, radicals like O, H, OH and HO2 form. When the combustion of hydrocarbons like cellulose and lignin in wood takes place, we get a greater variety of radicals, some of them carbon-centered. C2 and the radical CH arise in excited form and release blue and green light. Lignin-derived radicals involving benzene structures are not the healthiest things to inhale, but mere occasional exposure is probably nothing to worry about, unless there’s a concentration of wood-burning stoves in an area. All of this underlines the fact that whenever we write an overall equation for a fuel consumed in a fire, it’s like we are seeing only the ingredients of a recipe and the final product without witnessing the cooking.
Hot, gaseous products of combustion expand and rise, stretching flames vertically. The ascension leads to pressure gradients, and fresh air is pushed into the fire. The circulation supplies it with more oxygen, the electron-thief that campfires depend on to release heat as more tightly bonded products like water and carbon dioxide are created. There’s energy needed to drive molecular fragments of cellulose apart, in the same way that you need to exert force against gravity if you want to push a ball up a hill. But once at the top, the ball can roll further down on the other side. With chemicals, it’s not the combination of mass, gravity and varying heights that accounts for differences in potential energy but Coulombic forces acting over a variety of distances between positive atomic nuclei and valence electrons.
Why is a wood flame predominantly yellow orange? It has been proposed that it’s not the result of electron transitions; what’s responsible is the incandescence of particles at about 1100 to 1200 oC. Since the combustion of wood is incomplete, the flame’s soot particles, some of which are elemental carbon (others are polymers), emit part of their vibrational energy as photons. How fast the molecules vibrate depends on their temperature, and the hotter the surface, the higher the frequency of the photons emitted. The same mechanism accounts for the red glow of logs at the base of a fire. But the temperature is lower, about 700 to 1000oC, hence a color of a longer wavelength and a lower frequency. Different parts of the charcoal emit light of slightly different frequencies, intermittently and in different directions. A point on the surface of the charcoal particle that has just emitted photons will have lost energy and cooled slightly. Although exothermic reactions quickly compensate, from that same spot, the temperature will not necessarily be identical, especially considering air movement and the exact frequency of photons is not necessarily replicated.
It’s well known that the ease of ignition and burning rate of wood vary greatly with moisture content. Specifically, a 10% drop in moisture-content results in an increase of 20–30% in burning rate. When wood is too dry the combustion rate increases, but an inadequate oxygen supply leads to more undesirable emissions. The combustion rate also depends on boundary conditions and the species being burnt. Why does it vary with tree type? Wood composition is not constant. Wood is essentially a matrix of cellulose and other carbohydrate fibers (hemi cellulose) reinforced by the adhesive binding action of lignin. But hardwoods can have anywhere from 18 to 25% lignin along with varying amounts of hemicellulose, usually a partly acetylated, acidic xylan. Softer woods have other hemicellulose fibers and more “binder”, 25 to 35% lignin. There is also an assortment of oils and secondary products present.
The different wood recipes not only affect kinetics but thermochemistry. Softwoods, compared to hardwoods, release on average an extra 5% of heat, a maximum of 21 instead of 20 MJ/kg, to be precise. From the point of view of carbon dioxide emissions, it’s not a good idea to rely on wood as a primary fuel. For every MJ of heat obtained from wood, on average, 80 g of CO2 are emitted. In contrast, natural gas combustion only puts out 50 g per MJ.
ii. A hands-on example of oxygen’s explosive reactivity
When we were preteens, someone in our neighborhood showed us how to make a tennis-ball cannon. It consisted of cutting off the top and bottom of about 6 cans of beer and then taping them together. For the seventh can—the cannon’s base and reaction chamber—we left the bottom intact and cut off only a circular part of the can’s top so it could act as a support for a tennis ball. About an inch from the bottom we also poked about a 1 cm- hole and then taped the base to the rest of the can assembly. Through the hole, we poured in a bit of gasoline, slid the ball in, shook the cannon, waited and then ignited it with a match.
The first time we fired the cannon, we were in the fields behind our street. Its range astonished us as the ball almost inadvertently struck a woman on a 2nd floor balcony, hundreds of feet away. She was putting out her laundry, and she hollered at us when the ball bounced off the brick wall behind her. A week later, just before we tried it again, an inquisitive adult by the name of Eduardo laughed in disbelief when we told him what we expected to happen. The cannon was almost in a vertical position when we fired it. Eduardo’s jaw dropped almost as low as his clavicle. The ball had gone so high that we momentarily lost sight of it.
After hearing the story, my students were exploding with enthusiasm to build their own. In the early years of my teaching career, beer can-diameters had already been made incompatible with the size of tennis balls. But the biggest problem was that it would have been irresponsible to use gasoline around students. The heat of combustion for its safer substitute, ethanol, was no match for that of gasoline. Launching it became a hit and miss affair, and we did not initially get the bang, or the range promised by my recollection.
Understanding the theory helped us out in the long run. Before ignition, some alcohol molecules are in the vapour phase. Along with the air molecules in the base of the cannon, they exert pressure on the ball. But that pressure is not significantly different from the air pressure exerted on the other side of the ball, facing the cannon barrel’s exit.
After igniting the alcohol, we have a different matter. If there’s sufficient oxygen, the alcohol can optimally react with oxygen gas in a 1 to 3 ratio to yield carbon dioxide and water in a 2 to 3 ratio. The total potential energy of the oxygen and alcohol lie mostly within the bond energies of oxygen and alcohol’s intramolecular bonds. That sum exceeds that of the intramolecular bonds of the products of combustion, which is why the excess energy is released.
The sudden release of heat excites the molecules of the CO2, H2O and of the part of the air that did not react. As these molecules pick up speed in a confined volume, they exert far more force per unit area on the ball. The pressure deforms the ball, and as it regains its shape, some of the energy becomes the ball’s kinetic energy, accelerating it out of the cannon. But the gas-ball collisions are not elastic. Some of the energy bunches up air molecules, forming waves. In other words, some of the energy ends up as a whistling or thunderous sound, depending on how much heat was released. With an unsealed ignition-hole and reliance on tape, some exhaust gases escape, wasting even more energy.
One year I was in my brother-in-law’s steel workshop when I noticed some spare metal cylinders. I got him to weld one to a base, told him where I wanted the hole, and suddenly we had a much more solid tennis ball cannon. When I brought it to school, to get more of the alcohol vapours to ignite, to approach a more stoichiometric ratio, we decided to use a portable oxygen tank to inject gas into the chamber. To prevent it from escaping, we sealed the hole with plasticine through which we had inserted a firecracker fuse.
The firing of the cannon was spectacular(Figure 3), judging from the exploding sound and voices after the fuse went off. The ball zoomed across the parking lot. No one saw the ball in flight, but we found it because some of us heard it hit a tree that intercepted its trajectory. Here it is in action:
There’s more than enough muzzle velocity for the ball to detach someone’s retina, so safety goggles are a must. But that danger aside, it would be nice if gun users would some day replace their bullets and guns with alcohol-cannon projects.
Ozone(O3)
In various classroom demonstrations,I often used a Tesla coil. The one we owned looked like a giant pencil. Its removable tip was pointy and metallic. When plugged into an alternating current source, it would generate a high voltage and low current that could dimly illuminate an unconnected fluorescent tube without making contact. In between the tip of the coil and the pair of the bulb’s contacts, there would be a blueish arc, resembling a mini bolt of lightning. There would also be a faint, unpleasant chlorine-like smell that students sitting in the front row would detect. What was going on?
The blue color is already an indication that air molecules are being excited by the high voltage. But the smell results from the formation of ozone, O3 . The coil has enough energy to split each oxygen molecule into a pair of oxygen atoms, one of which then bonds to an unsplit molecule of O2 to yield O3.
The commercial generation of ozone is always done on site because the molecule is unstable. A high voltage alternating current (600 to 20 000 volts) is applied across a dielectric discharge gap that contains either air or pure oxygen (Figure 4). (A dielectric is a ceramic material that keeps the charges separated.) Between the plates we observe a bluish corona, like the plasma we saw at the tip of Tesla coil when it approached a metallic surface, but on a larger scale.
At a water treatment plant, ozone is a friend. It is better than chlorine at killing harmful microorganisms without leaving chlorinated hydrocarbons behind. Ozone, unlike chlorine, is even effective at breaking down most pesticides. An ozone dose of 0.4 mg/L for only 4 minutes is effective for pre-treated water. In 2016, the city of Montreal treated about 800 liters of water per person daily, of which about 75% was treated with ozone. Although the city’s consumption habits have been improving with better designed flush toilets, less watering of lawns and replacement of old water mains, they are still wasteful compared to most European cities.
Ozone unfortunately breaks down faster than chlorine, so some of the latter is still used, in case drinking water is recontaminated between its source and destination. But use of ozone significantly lowers the concentration of chlorine needed, rendering the water healthier and almost odorless. Many cities also treat sewage with ozone in the final stages before releasing it into the rivers to reduce bacteria and simplify the soup of compounds that would otherwise be ingested by aquatic life and people downstream.
In more recent years there have been many commercial attempts at retarding fungal growth on transported fruits using very low levels of ozone. A 2016 review study concluded that it’s not clear how efficient the gas really is at reducing spoilage. In medicine an ozonized oil seems to be effective against dermatophytes that attack soles of the feet.
Outside of a controlled environment, ozone is undesirable at ground level. Like the chlorine whose smell it resembles, it is a strong oxidizing agent, attacking rubber, eyes and lungs. Once oxygen was introduced into atmosphere by photosynthesis, another event of life-induced chemical evolution occurred. Up in the stratosphere some of the ultraviolet light from the sun, UVC, acted like the Tesla coil, creating ozone from oxygen. Molecular oxygen continues to protect us from dangerous UVC, whose wavelengths are in the 100 to 200 nanometer regions(Figure 5).
Figure 5. Ultraviolet absorption of molecular oxygen
However, as Figure 5 indicates, molecular oxygen does not absorb well in the 200 to 280 nm region(UVB) of the ultraviolet spectrum. Ozone, instead, is capable of absorbing in that ultraviolet region (Figure 6). When each ozone molecule takes in energy of that wavelength it becomes one diatomic oxygen and one oxygen atom. Heat is released, keeping the stratosphere warmer than the troposphere underneath it but also providing atomic oxygen to keep the ozone layer at steady state.
Of course, we all know, thanks to the efforts of Rowland and Molina, that the steady state has been threatened, especially over Antarctica. Polar stratospheric clouds trap NO2 from volcanic emissions. That gas would normally neutralize the Cl from molecules like Freons which were common in refrigerating coolants. But once free, the Cl would compromise ozone concentrations by inducing its own steady state. How? Although the exact mechanism varies with latitude, basically the story goes like this. Cl is a neutral atom of chlorine with a valence of 7, a so-called radical because the 7th valence electron is unpaired. First the Cl radical plucks out an oxygen from O3 to create O2 and the radical ClO. Then ClO reacts with an ingredient of ozone-synthesis, atomic oxygen, to produce molecular oxygen and a Cl radical. The latter is now available to re-attack ozone, repeating the destructive cycle.
Hydroxide ion (OH–)
If you collect the white ashes from a wood fire, you will have a source of potassium carbonate, K2CO3. This water-soluble compound, when added to water, will create a solution that feels slippery between the thumb and the index. Why?
Water has an extremely small concentration of positive(H+)ions negative OH ions. These OH ions (OH–) called hydroxide have a negative one charge and without them we would not be able to dissociate water into oxygen gas. They are in equilibrium with unsplit water molecules and normally the positive and negative ions exist in equal numbers, keeping water’s pH neutral. While the negative-two carbonate ion uses up an H+ from the water to become hydrogen carbonate, the baking- soda ion that flows in blood and rivers, water is left with an excess of hydroxide. It has become a base or an alkaline solution. Some of those ions break up the oils of the skin into fatty acids, ingredients of soap—hence the slippery feel.
This is essentially how soap is manufactured on a large scale. Instead of skin oils they use vegetable fats like low-quality olive oil, palm and coconut oil. And instead of potassium carbonate, they use potassium hydroxide (KOH), which is a stronger base. Since the reaction also produces glycerin, salt is used to separate the two products. After removing the glycerin from the top, at this stage there is still some unreacted fat. An even more concentrated solution of hydroxide is added, and then additional steps are required to purify the soap. Finally, to make it more appealing, scents are added to the soap in one of the last stages of manufacturing. As one could see, it’s not just a one-shot deal that one imagines from seeing a chemical equation of the main reaction. The literal interpretation of the equation just leads to a mess in the lab!
To mention that potassium hydroxide is also used to make alkaline batteries, synthetic fertilizer and paper would be describing just the tip of the iceberg. It is a highly versatile compound. They once made it by reacting limewater with potassium carbonate. Nowadays, both sodium hydroxide (NaOH) and KOH are made by electrolyzing the respective aqueous chloride solutions. Chlorine and hydrogen gases are both byproducts of this reaction and are collected. Interestingly this is a side reaction if you try to electrolyze tap water, which has chlorides as impurities. You can always tell which students(or technicians and teachers!) are being honest in reporting their data. The expectation of a 2:1 ratio from the reaction of pure water skews many of the results. The side-reaction, in fact, creates a higher ratio because not only is it a second source of hydrogen, but more importantly, while the chloride ions compete with hydroxide to give up electrons, less oxygen gas is produced.
Rust (various oxides and hydroxides of iron)
I recall hearing about a woman who, decades ago, had a rust-prone Dodge model that was free of rust. It wasn’t because she kept the car locked up but because, unlike most people who buff their cars only in the summer when they need it the least, she waxed her Dodge a dozen times over the winter. In the overall reaction of atmospheric corrosion of iron, the reactants are iron, oxygen and water. Washing and drying prevents dirt from trapping moisture and washes out salt that accelerates corrosion. Waxing sets up a barrier between oxygen and iron.
The product of rust, contrary to what basic chemistry textbooks say, is more than just Fe2O3. The rust layer consists of a variety of minerals, including magnetite (mainly Fe3O4), maghemite(ϒ-Fe2O3), goethite, lepidocrocite and akaganeite ( mostly α, ϒ and β forms of FeOOH, respectively), reduced lepidocrocite (ϒ-FeOHOH) and ferrous hydroxide(Fe(OH)2). Such a minestrone is a hint that the chemistry is more complicated than previously imagined.
In the first stage of rusting, iron or low-alloy steel loses electrons and becomes soluble Fe2+. But the electrons are only partly taken up by oxygen. The majority are accepted by another ionic form of iron, Fe3+. Only after the latter are used up does oxygen pick up the slack. Along with water, oxygen picks up electrons to produce hydroxide which bond with Fe2+ to create a variety of iron compounds, which are then further oxidized. As the water from the electrolyte film evaporates, it slows down the whole process. Since humidity influences the rate of evaporation, it determines how many cycles of this process iron experiences over time.
Intermediate compounds with Fe3+ have low solubility. This slows down further formation of rust, but in the presence of sea salt or street salt, they are displaced by chloride ion. As Fe3+ ions go into solution, they are subject to a variety of reactions known as iron hydrolysis. Specifically, they react with water to form three different hydroxide complexes and, in each case, acidic ions. The acid then accelerates corrosion. Any acid already in the environment due to acidic precipitation will be even more efficient at producing rust.
There’s a connection between corrosion and the evolution of life. After the evolution of photosynthesis, almost all the O2 made by early cyanobacteria was used up in converting large amounts of Fe2+ in the oceans to Fe3+. This reaction precipitated prodigious amounts of ferric oxides, leading to banded iron formations in sedimentary rocks. More importantly, it led to a very slow accumulation of oxygen in the atmosphere. Oxygen was toxic to all forms of life, but the gradual increase gave time for cellular respiration to evolve in bacteria.
Hydroxyl radical (.OH )
If we somehow developed a ratio of a chemical’s importance to how well-known it is to a scientifically literate audience, the hydroxyl radical (.OH ) would probably win the prize. A popular freshman chemistry textbook in North America does not even list hydroxyl as a topic in its index. Yet when it comes to cleaning up after impurities in the atmosphere, it is the most important oxidizing species in the air.
Two molecules of the so-called detergent of the atmosphere are formed for every oxygen radical that reacts with a water molecule. When the pollutant nitrogen dioxide is removed from the atmosphere it is the hydroxyl radical that turns into nitric acid. Similarly, it converts sulfur dioxide pollution into sulfuric acid, although the formation of acids isn’t without consequences.
But there are very few hydroxyl radicals in the air for two reasons. Most oxygen radicals formed by ozone’s ultraviolet dissociation go on to form new ozone. Only 3% combine with water to form hydroxyl. The other reason is that they are short-lived due to their high reactivity, existing for only 0.01 to 1 second after they are formed. This makes hydroxyl radicals extremely difficult to measure. Their tracking must be indirect, usually based on measuring atmospheric chemicals whose sources are well-known and initiated mostly by a reaction with . Even when their concentration is boosted by the tropics’ humidity, more intense sunlight and higher temperatures, their concentration rarely exceeds 2 parts per trillion(ppt) at sea level. Elsewhere on Earth it can be as low as 0.01 ppt (Figure 7).
Figure 7 Hydroxyl radical concentrations across the globe. The unit of 10 million radicals/cm3 is equivalent to 0.5 parts per trillion(ppt) at sea level pressure, so a 3.0 on the colored scale is 1.5 ppt.
When we release methane into the atmosphere, hydroxyl radicals turn it into water and a methyl radical which goes on to form formaldehyde. Unfortunately, every molecule of methane, a greenhouse gas, remains in the atmosphere for about 10 years before it’s cleansed by hydroxyl. Luckily, the poison carbon monoxide’s residence time is only 3 months. But upon attack by hydroxyl, it becomes the greenhouse gas carbon dioxide.
In short, we must be thankful for hydroxyl’s existence. The air quality would certainly be worse without it. But because of its slow kinetics and some undesirable products, it is not the solution to our pollution. When carbon monoxide is present, it consumes about 70% of hydroxyl in its immediate environment. Much of the other 30% attacks methane, which regenerates carbon monoxide. In Chemistry of Atmospheres, Richard Wayne pointed out that since an average adult male ingests a total of 3.2 kg of food and water but inhales 13.5 kg of air per day, the air quality is at least as important as the cleanliness of food and water.This includes indoor air quality, and unfortunately hydroxyl’s role in that context has received far less attention. Due to the radical’s short lifetime, it cannot be assumed to be transferred from outside air. But it does exist inside homes. In 2013, Gligorovski directly measured concentrations of 1.8 × 106molecules per cm3. It forms when short wavelengths of light breakdown HONO, a gas derived from NO2, which in turn forms from cigarette smoke, fireplaces, furnaces, car exhaust (from outside or a garage) and from gas stoves.
Advanced oxidation processes(AOP) are being developed to treat municipal and industrial waste take. These techniques take advantage of hydroxyl radicals. They are usually generated by coupled chemical and/or physical systems including combinations of hydrogen peroxide (H2O2) with ferrous ion (Fe2+)ferric ion (Fe3+) catalyst and ozone. To increase the concentration of the radical, ultraviolet lamps are used. Analyzing the intermediate reactions and products is a challenging task. Many occur in less than 100 microseconds. Research into AOP’s chemistry, not the technique itself, relies on a combination of pulse radiolysis (which uses highly accelerated electrons from a linear accelerator) and competition kinetics methods, laser-induced fluorescence, and electron spin resonance (ESR). The latter resembles proton nuclear magnetic resonance(1H NMR), but it excites the spins of electron instead of those of protons. ESR spectroscopy (Figure 8) was developed to study radicals such as hydroxyl. Medical research also makes use of ESR to examine the link between some diseases and oxidative damage by radicals such as .OH.
Figure 8. Three ESR spectra. The squiggles in the top graph are those belonging to hydroxyl. Its peaks are no longer visible in the other two spectra because they have been scavenged by metallothionein isoforms. The radicals were generated by Fenton reactions, which are also used in advanced oxidation processes (AOP) discussed earlier. Source: Proceedings of the Western Pharmacology Society 41:155-8
The face of sulfur is like that of chemistry. It finds its way into compounds that are poisonous and/or foul-smelling to us, making it easy for us to conjure a negative but starkly incomplete image of sulfur. In the bigger picture, however, there would be no life and less rain without its compounds. If sulfur existed, but if we had never learned how to modify the elemental form that occurs naturally, our world would be radically different.
Orthorhombic sulfur (S8) is the most common of several possible, so-called allotropes, molecular combinations involving a single element. It is one of the few elements known to the ancients because it can occur unbonded to other elements near volcanoes and hot springs. Egyptians used it on wheat 2000 years ago to prevent rust disease.
The element is bright yellow, and if you handle it to feel its smooth texture, you will walk away with smelly hands. When I was young, a kid in our neighbourhood, François, invited me to take a whiff of a sample of burning sulfur. (With a match, he had lit a sample from his chemistry set.) It turned to an attractive blue colour. The unusual flame incited me to move closer, but the smell felt like needles going up my nostrils. Of course, François laughed at my reaction. When I saw it again months later in my brother’s chemistry set and in school, I made sure not to burn it with my nose nearby; yet the indelible mark it left did not overcome the appeal of its distinguishing color.
i. Sulfur from tar sands
Since childhood I have encountered elemental sulfur countless times, even using it to clean up spills of mercury when it was still allowed in high school labs. Not once did I wonder about its origins until recently when I saw hills of sulfur at the port in Vancouver(Figure 1). It has apparently been a common sight for a long time, and it comes from Alberta’s tar sands.
Figure 1. Elemental sulfur at Vancouver, Canada’s port. It likely comes from Alberta’s tar sands. Picture by author.
There is a fair amount of sulfur in their heavy crude. In the refining process poisonous hydrogen sulfide, H2S, is produced, which is absorbed and separated from hydrocarbon gases by an alkaline (basic)compound. The sulfur is finally extracted by the Claus Process, which consists of a two-step-process: high- temperature oxidation of H2S and conversion to sulfur with a catalyst. A lot of sulfur is extracted this way, more than can be sold, and with the excess they decided to build pyramids out of pure sulfur, just outside of Fort McMurray, Alberta (Figure 2).
The process does not remove all sulfur, which becomes a problem because burning heavy crude produces not only carbon dioxide but sulfur dioxide. From 2018 onwards, global regulators were trying to crack down on sulfur dioxide emissions, which could shrink the market for heavy crude and force Canadian oil sands companies to further discount their product.
What happens to sulfur that’s sold and not used to make pyramids? Its most common destiny is sulfuric acid, to be explored in the next section. Sulfur can also be used to make matches, dandruff shampoo and fertilizer. In the elemental state, it is used as a fungicide in organic or industrial farming. Although there are “sulfur-shy” plants such as gooseberries, apricots and raspberries that should never be treated with sulfur, others tolerate the element, which can prevent powdery mildew, rose black spot and fungal rust. Since sulfur is a preventive fungicide it will not work if the diseases have already appeared.
ii.Sulfur from photosynthetic bacteria
In 1931, sulfur gave an open-minded Dutch-American graduate student, Cornelis B. van Niel and eventually the rest of the rest of the world new insight into photosynthesis. Purple sulfur bacteria found in salt marshes and mudflats produce organic matter in the presence of light, but they rely on hydrogen sulfide instead of water. The waste-product is not oxygen but sulfur.
Figure 3. Chromatium species, a type of purple sulfur bacteria from a salt marsh pool near Woods Hole Oceanographic Institute, Massachusetts. This microorganism produces sulfur instead of oxygen when photosynthesizing. Source: http://cfb.unh.edu/
For the previous 100 years, scientists were under the false impression that in photosynthesis the entire water molecule was incorporated into sugars. They also incorrectly imagined that water effectively displaced oxygen from carbon dioxide. Seeing that purple sulfur bacteria(Figure 3) displaced sulfur from H2S, an analogue of H2O, Cornelius B. van Niel, correctly inferred that in plants, oxygen was displaced from water molecules.
Eventually more solid evidence supported van Niel’s hypothesis. Plants were given water that was made with the 18O isotope, a non-radioactive but heavier version of the common 16O isotope. They were grown in the atmosphere with the usual CO2, which has mostly 16O. When the oxygen produced by the plants was analyzed, it contained only 18O, proving that water is the source of oxygen gas.
iii. Sulfur from Io
Figure 4. Io with its sulfur(S8) surface. The small areas of orange-red are S3 and S4 forms. Source: NASA
The Hubble telescope was equipped with the Space Telescope Imaging Spectrograph (STIS) which revealed information about an object’s chemical content, temperature, density, and motion. Jupiter’s moon, Io, as any colored photograph hints(Figure 4), is covered with a thin coat of sulfur, which is emitted by its volcanoes. In 1999, STIS detected S2, a diatomic version of sulfur, above the volcano, Pele. S2 lands on Io’s cold surface, where sulfur atoms rearrange into tri- or tetra-atomic molecules (S3, S4). These lend a reddish color to the surface. Eventually the atoms rearrange into their most stable configuration, molecular rings of eight atoms (S8). This is the form that we see along the edges of cauldrons in Yellowstone Park. Interestingly, in the extraction of sulfur from heavy crude, sulfur is also initially in an unstable S2 form but then becomes S8, the form piling in the Fort McMurray pyramids and in the piles at the Vancouver port.
Sulfuric acid
In concentrated form, sulfuric acid, H2SO4, is a thick syrupy liquid. I recall using it while working in a quality control lab analysing the protein content of hot dogs. Despite the pressure differences of the fume hood which are meant to divert the acidic fumes away from the working environment, one of my colleagues became itchy all over her arms and torso. She reported to the company doctor but regretted her decision when she felt that he had taken advantage of her sensitive reaction to the acidic fumes. He had asked her to disrobe and expose her breasts. It was the first time in my young life that I had heard of unprofessional behaviour from an esteemed individual.
Industry’s most widely produced compound, H2SO4, is used in lead batteries, copper refining and especially in the production of phosphoric acid for phosphate fertilizers. It’s also needed to make other acids, pharmaceuticals and dyes. Currently, sulfuric acid is made industrially by burning sulfur and by catalytically oxidising the subsequent product (SO2) with vanadium(V) oxide. Contrary to many internet sources, the new product, SO3, is not added to water in the final step. That reaction is too exothermic for a large-scale operation. Instead SO3 reacts with previously made sulfuric acid to make H2S2O7, which then reacts with water to yield the acid.
When SO2 escapes into the environment, it becomes sulfuric acid through the reaction of hydroxyl radicals and moisture. This main component of acid rain has a serious impact on lungs, infrastructure and lakes. For this reason, industries equipped with the latest technology have scrubbers. Some of these use limewater to convert the SO2 into calcium sulfite and eventually into the useful CaSO4 · 2H2O, which is the material in wallboard.
If you are going to be a criminal, you don’t want to be a popular one or you will increase the likelihood of being held responsible for crimes that you did not commit. A great deal has been written about why onions make you cry, but the information contradicts. Is sulfuric acid being falsely “accused”? And how we’ll arrive at the truth might prove to be more fun than knowing the fact itself.
The 2012 chemistry textbook, General, Organic and Biological Chemistry, by H. Stephen Stokes; a PBS website on why onions make us cry, and chemistry.about.com all claim that, after the volatile lachrymator propanethial-S-oxide is produced by a cut onion’s enzymes, it reacts with water in the eyes to produce the irritant sulfuric acid. But Scientific American, Molecule of the Month and the analytical services of onionlabs.com attribute the lacrimatory effects directly to the molecule’s action on receptors. Sulfuric acid is not part of the explanation. From the latter:
“LF is the chemical compound which directly causes the eye to tear (often called the onion lachrymator) and the chemical sensation of heat or mouth burn when an onion is eaten. It is measured using HPGC equipment and is reported in µmoles of LF/ml of onion juice. . ..“
And then the details from a researcher writing in Scientific American:
“The cornea is densely populated with sensory fibers of the ciliary nerve, a branch of the massive trigeminal nerve that brings touch, temperature and pain sensations from the face and front of the head. The cornea also receives a smaller number of autonomic motor fibers that activate the lachrymal (tear) glands. Free nerve endings detect syn-propanethial-S-oxideon the cornea and drive activity in the ciliary nerve–which the central nervous system interprets as a burning sensation–in proportion to the compound’s concentration. This nerve activity reflexively activates the autonomic fibers, which then carry a signal back to the eye ordering the lachrymal glands to wash the irritant away.“
I started to get skeptical about the alleged formation of sulfuric acid for the following reason. If, in the presence of water in the eye, the lachrymator degrades so quickly, wouldn’t most of the propanethial-S-oxide break down in the onion’s aqueous environment? Let’s assume for a moment that some still escapes unscathed due to its relatively high vapor pressure of 41.2 ± 0.2 mm Hg at 25°C. To put that in perspective, water’s vapor pressure is only 23.8 mm Hg at that temperature, so the lachrymator is about as volatile as ethyl alcohol, whose vapor pressure is about 44 mm Hg.
To look for any pH-changes, I taped only the ends of a wet piece of litmus paper in a sealed plastic bag (Figure 5). I chopped an onion and quickly placed the pieces into a Ziploc, making sure that the acidic juices of the onion would not be in contact with the litmus. If the lachrymator indeed broke down into compounds that included sulfuric acid, then the exposed wet litmus should turn red between 30 seconds and couple of minutes, the time that it takes for the compound to form and reach the eyes. But there was no color change, consistent with the fact that the lachrymator itself has no pKa, a measure of acidic strength.
One could argue that the amount of acid produced was too small to affect the litmus. A paper on a rapid extraction method of the lachrymatory factor of onion found that each milliliter of onion juice contains 1−22 μmol of the tear-forming agent. Let’s assume a density of 1.0 g/ml and an average mass of 120 g per onion, which is about 90% water, and that about only 30% of the lachrymator reached half a ml of water on the litmus. If it had produced H2SO4 in a 1:1 ratio (only one source (the textbook) had included a chemical equation, and it was not even balanced), then we could easily calculate that we’d end up with a pH between 3 and 4, enough to turn the litmus red. I also mixed a small amount of vinegar with the onions, and within a few minutes, the same wet litmus did indeed turn red, revealing that the simple setup could detect acidity.
And one final detail, most literature on the topic including this blog so far, discuss the lachrymator as if it’s a single compound. But an NMR analysis in Tetrathedron Letters reveals it to be a 19 to 1 mixture of (Z)- and (E)-propanethial S-oxide (Figure 6).
Figure 6. Z-propanethial-S-oxide makes up 95% of the lachrymator in onions.
These are not earth-shattering revelations. But what was humbling is how the inaccuracies from a field that I know and enjoy almost got past me. Imagine what happens when I read news about economics and terrorism, stories that I often do not dig into to uncover the truth.
Sulfate, the essential ion
Sulfate, SO42-, is a part of the sulfuric acid molecule, but if it’s bonded to a metal, the compound will not share acid’s most distinguishing properties. The sulfate ion itself does not burn skin or taste bitter. SO42- is also part of the planet’s natural sulfur cycle. Algae and plants absorb sulfate ion to make a pair of amino acids that contain sulfur, cysteine and methionine (Figure 7).
Figure 7. Sulfur-containing amino acids.
Humans cannot make their own methionine. Obtaining it from our diets, we ultimately rely on other organisms’ ability to do something with sulfate. Even when we make cysteine, we get the sulfur from methionine. The sulfur linkages between different cysteine molecules in a peptide chain help it fold and stabilise into a specific three-dimensional protein that define its role. The importance of shape and role can range from how curly your hair is to the proper functioning of enzymes in our cells, and without methionine no nucleated cell can even begin to make a protein. Bacteria in ruminants also use sulfate to produce cystine and methionine. This was discovered by using sulfate with a radioactively tagged isotope of sulfur.
Elsewhere sulfate is reduced to H2S by specialized bacteria while others can oxidize either H2S or elemental sulfur back to sulfate. Bear in mind that H2S is also produced by volcanoes along with sulfur dioxide, which can react in the atmosphere to yield sulfates again. Sulfate-containing particles are important in cloud formation; they serve as seeds for the condensation of water vapour.
If as a teacher or student you hear of an oxidation-reduction partnership, the words might conjure up an image of people partnering up to do a redox lab, which involves one substance losing electrons and the other sweeping them up. But until recently, not even the most specialized biochemists imagined that two different organisms were symbiotically working together to oxidize methane and reduce sulfate, respectively, in layers of sediment and rock under the sea floor. One of the microbes, a methanotroph which uses methane (CH4) as its energy and carbon source, reduces CH4 to hydrogen carbonate ion (HCO3–) in the absence of oxygen(Figure 8).
Figure 8. Two microorganism cooperate to perform an oxidation reduction- reaction.
But the metal ions that the organism uses to pick up electrons lost by methane are not efficient enough. A sulfate-reducing bacteria comes to the rescue. It uses the electrons released by the oxidation of methane to reduce sulfate (SO42-) to sulfide (S2-). Its reward? It too obtains energy in the process.
Sodium sulfate is found in three forms: (1) Glauber’s salt (Figure 9) which is a hydrate of the salt(Na2SO4.10H2O); salt cake (mostly Glauber’s salt but mixed with magnesium sulfate and silica;and finally anhydrous sodium sulfate(Na2SO4). Sodium sulfate is needed to make brown paper, brown cardboard, glass, textiles and detergents. In 2016 China produced 60% of the world’s sodium sulfate and consumed 67% of the world’s share. Chaplin Lake in Saskatchewan, Canada has large natural reserves of salt cake, which is converted to 99.0% anhydrous sulfate, the so-called detergent-standard. To obtain that purity, the salt cake has to go through a log washer and then through a heating process that drives off the water from the hydrate.
Figure 9. The mineral mirabilite is Glauber’s salt.
Sulfur dioxide
SO2is a useful gas in the production of sulfuric acid, but it is also an undesirable combustion product of coal, which is typically 1 – 4% sulfur. Nickel and copper ores are sulfur compounds, so in the roasting process, sulfur dioxide is again produced. Scrubbers are essential in filtering emissions and converting the smelly and lung-irritating gas to useful products. These were discussed in the sulfuric acid blurb.
Emission control systems are hindered by the presence of sulfur dioxide in exhaust. The catalyst in converters is designed to oxidize some of the carbon monoxide and dissociate some nitrogen oxides. SO2 remains intact and interferes with the catalyst. As a result, there has been pressure to remove more sulfur from transportation fuels. If the crude oil is originally sulfur-rich, it is dubbed “sour. Sweet crude oil contains less sulfur, and its combustion poisons less catalyst in converters and contributes less sulfur dioxide to the atmosphere.
Volcanoes are a natural source of sulfur dioxide. I once hiked near Kīlauea volcano in Hawaii and the SO2-containing volcanic smog was foul-smelling and made it difficult to breathe on the steeper parts of the hike. During major eruptions, SO2 can be injected to an altitude of over 10 km and into the stratosphere. There, SO2 is converted to sulfate aerosols which reflect sunlight and have a temporary cooling effect on the Earth’s climate. They also deplete ozone.
The overall volcanic contribution of SO2 to our atmosphere, however, pales in comparison to the amount we discharge. Emissions in North America and Europe peaked in the mid-70s but are currently at and of their maximum values, respectively. But meanwhile in east Asia, the emissions have reached Europe’s highest output from 40 years ago.
Sodium bisulfite (NaHSO3)
Sodium bisulfite, a source of the bisulfite ion, is made by reacting washing soda (sodium carbonate) with sulfur dioxide, the same starting point in the contact process-production of sulfuric acid. Water treatment plants use bisulfite to eliminate excess chlorine from both drinking water and treated sewage. Relative to chlorine, bisulfite acts as a reducing agent, allowing chlorine to gain electrons and become chloride. In a similar reaction, bisulfite can react with oxygen to prevent corrosion of large pipes.
An oxidation-reduction reaction involving bisulfite has a strong, peculiar smell that you can feel in your throat because the bisulfite ion and an acidic proton are in equilibrium with water and sulfur dioxide. Amateur wine makers use bisulfite as a reducing agent to prevent the oxidation and discoloration of white wine. They must be careful with concentrations; otherwise too much will leave a sulfurous taste that ruins the wine. Even when professionals are careful, they run up against individuals who are allergic to bisulfite, which makes them wheeze and develop hives.
Hydrogen sulfide (H2S)
The odour threshold for hydrogen sulfide gas is only 10 parts per billion to about 1.5 parts per million. It stinks like rotten eggs or like the nastiest type of flatulence, both of which involve the microbial breakdown of sulfurous amino acids into hydrogen sulfide. Luckily, we are sensitive to the smell because at 1 to 2 parts per thousand it causes instant death. In between the odor threshold and the lethal dose, it has unpleasant effects ranging from nausea, pulmonary edema and loss of consciousness. Since hydrogen sulfide is produced naturally from decaying organic matter, it can be released from sewage sludge and manure. Sulfur hot springs and natural gas are also sources. Some industrial processes that release it as a byproduct include wastewater treatment, petroleum refining and mining. Workers have died when ventilation and detection systems have been inadequate. Marriages have been threatened by the indiscretions of a partner suffering from H2S flatulence.
Let us not forget that one man’s poison is another organism’s lifeblood. Without H2S, purple sulfur bacteria would not be able to make sugars. It provides electrons and hydrogen ions that sustain their form of photosynthesis. In the South Andros Black Hole, Bahamas, a thick layer ofpurple sulfur bacteria lives at a depth of 17.8 m. In this thick layer which makes the water appear black at the surface, the temperature increases sharply to 36 oC as a result of the inefficient transfer of light from their accessory pigments, carotenoids, to their chlorophyll centers. This ability to raise the temperature of their immediate environment gives the mass of bacteria an advantage over other species that are incapable of the feat. And it’s all dependent on H2S.
In both the chloroplasts and cytoplasm of phytoplankton, the sulfur-containing amino acid, methionine, is converted into dimethyl sulfonium propionate (DMSP). Why? DMSP is an organic salt that prevents water from leaving microscopic algal cells that live in the ocean’s salty environment, and at the same time DMSP does not interfere with protein function. The compound also helps bacteria that consume dead plankton.
Eventually, DMSP is broken down to acrylic acid and a vital component of the sulfur cycle known as dimethyl sulfide (DMS). It has been suggested that DMS and another compound known as isoprene are used by phytoplankton to mitigate harmful reactive oxygen species like H2O2 formed from dissolved oxygen, organic matter and iron stress.
Annually, approximately 28 × 1012 grams of sulfur are placed in the atmosphere when algae release the volatile DMS into the atmosphere. The compound is eventually oxidized to sulfate, which is an important cloud condensation nucleus. Although it hasn’t been substantiated by evidence, Charlson proposed a mechanism involving DMS that became part of Lovelock’s Gaia hypothesis. In this scenario, the abundance of cloud condensation nuclei from DMS-derived sulfate above the sea affects the amount of cloud cover, in turn affecting the radiation reaching the ocean. This feeds back into the climate, affecting marine biology and presumably regulating the amount of DMS produced by phytoplankton.
Sulfur compounds in minerals
This rock contains a pair of minerals, both of them sulfides: (1)the golden-colored pyrite a.k.a. fool’s gold ( iron sulfide), and sitting underneath is (2) sphalerite, which contains zinc sulfide the main ore of zinc. Galena (PbS, not the TV network but lead sulfide) is also found often together with sphalerite.
Left, Elemental sulfur on aragonite , a mineral consisting of calcium carbonate. From Sicily. Right, Stibnite, which is mostly antimony sulfide, Sb2S3, is the principal ore of antimony.
People like to include precious stones in jewelry. It is obviously not just about the stone; there is workmanship involved in adorning the stone; there is the association with status and enhancing the interest, if not beauty, of the person who owns and wears the mineral. But there is something more beautiful about minerals in their own natural state. Here is a look at those containing species of sulfur.
Gypsum contains calcium sulfate dihydrate, CaSO4·2H2O, the same mineral that is found in wallboard. As is the case with many minerals, gypsum’s color can vary depending on the impurities it contains.
The colorless crystals consist of a sulfate of barium, known as barite. BaSO4 has a low solubility, preventing the otherwise poisonous barium from being released. The compound is opaque to x rays and in suspension, it is used in medicine for imaging the gastrointestinal track. Barium sulfate is also used in fireworks to create a green color. Source of sulfate mineral images: Wikipedia commons.
Sciences in the Mural of Life 