Sunday, November 11, 2018


About 150 million years ago, some living creatures started evolving ways to make their own light, and this is what bioluminescence is: The ability of an organism to create and emit light. Having your own "flashlight" happens to be very useful. Depending on the creature, it is used for mimicry, for camouflage, to warn away predators, for communication among one's species and to find mate.

Just as bioluminescence has a multitude of functions, it appears to have evolved in a multitude of different ways. In bioluminescent bacteria, the bioluminescent machinery appears to be borrowed from its machinery for cellular respiration. In marine organisms, it seems to have evolved from once-essential cellular detoxification machinery. In at least one dinoflagellate species, the photosynthesis mechanism has been fine-tuned toward light production. In the firefly, molecules that once broke down fatty acids for energy storage now emit that energy as light instead. In this article we will explore the mystery of how bioluminescence evolved and how it offers important life advantages of non-glowing relatives.

What Is Bioluminescence (And What Isn't)?

Although bioluminescence is quite rare if you measure it as a percentage of the total number of all of Earth's species, it is quite common among marine species and it is surprisingly diverse among different distantly related groups of organisms, including land organisms. Most bioluminescent species (about 70%) are marine and these include ocean-living bacteria, dinoflagellates (these are the tiny creatures that make disturbed ocean water sparkle with tiny green-blue lights), marine crustaceans such as some krill and shrimp, some echinoderms such as sea stars, sea lilies and sea cucumbers, as well as sea squirts, and a great many species of fish, and even a few sharks. Bioluminescence has also found its expression on land but it is not as common. Land species include some bacteria and fungi as well as some insects, annelids (worms) and arachnids (spiders, scorpions, ticks and mites etc.). Groups that do not contain any bioluminescent species are the land vertebrate classes such as amphibians, reptiles, birds and mammals to which we belong.

There are, however, several notable fluorescent land animals, such as the polka dot tree frog, a reddish yellow frog that glows bright green under ultraviolet light. Biofluorescence is also seen in many fish and corals, in jellyfish, butterflies, parrots, spiders and even in the flowers of the common four o'clock plant (Mirabilis jalapa). Unlike bioluminescent species, which make their own light, these organisms never glow in complete darkness. They absorb light and then emit it immediately once it is absorbed, usually at a longer (lower energy) wavelength. It is important to distinguish between a fluorescent species and a bioluminescent species. They are easily mixed up.



All scorpions such as this (actually black) one fluoresce bright aqua blue under an ultraviolet lamp. They glow more faintly in nature, in the dim ultraviolet light reflected by the moon. Their exquisitely sensitive eyes also happen to see this particular colour best.

Biologist Edie Widder offers us a taste of bioluminescence across the marine world with some incredible footage of bioluminescent marine creatures in this 13-minute video, part of her TED talk in 2013:


art farmer:Wikipedia

Common Eastern Firefly (Photinus pyralis)

This firefly species is common in North America. At twilight, males use flashes of greenish-yellow light to attract females, who will respond with an answering flash of their own.


Wofl-commonswiki; Wikipedia

Female glowworm (Lampyris noctiluca)

This glowworm is actually a beetle because it has a hard shell or carapace. In this case it is the much larger wingless female rather than the male that uses light to attract males. The brighter her glow, the more fertile she is.

In addition to these well-known land examples, bioluminescence has found widespread use among marine species that utilize it in ocean water. They use it near the water surface on dark moonless nights or very deep in the ocean, where sunlight can't penetrate.

NOT BIOLUMINESCENT (trick example):

Sierra Blakely;Wikipedia

Crystal Jelly (Aequorea victoria)

This jellyfish is commonly found floating and swimming off the west coast of Canada and the northern United States, especially in Puget Sound. This photograph shows you what the colourless animal looks like but you are actually seeing light reflected from the camera, a common misconception in such photographs found online. Its bioluminescence is only visible in its outer ring, as a faint blue-green glow (which is emitted only when it is disturbed), shown below.

Photo taken by Osamu Shimomura

Dr. Shimomura, a famous organic chemist and marine biologist, isolated two luminescent proteins from this species of jellyfish. One of them is a bioluminescent protein called aequorin. It emits blue light. He also found trace amounts of another protein, and this one is fluorescent rather than bioluminescent. It's called green fluorescent protein (GFP). This protein absorbs the blue light emitted from the aequorin complex and re-emits it as green light. The aequorin complex mentioned here is a substrate/enzyme complex - more specifically a luciferin/luciferase complex. We will get better acquainted with these two evocative-sounding words later on. GFP is now widely used in biological and medical research. It is this discovery for which Dr. Shimomura received the Nobel Prize in Chemistry in 2008.

Exactly what does this jellyfish uses bioluminescence for? That remains a mystery. Perhaps it's a warning to potential predators. Researchers know that individuals do not flash at each other and they do not glow continuously. They can be stimulated to glow when they are disturbed but it is rarely observed in undisturbed individuals.

Marine bioluminescence is very prevalent among those species that live in very deep pitch-black ocean water were no sunlight penetrates, a kilometre or more beneath the ocean surface. One could imagine that these animals would have lost their sight over time, like blind cave-dwelling salamanders did. Many of them have, relying on pressure changes and smell instead. However, some species evolved extreme light sensitivity to the faint bioluminescent light shows they encounter. Down here, water is cold, pitch-black and under tremendous pressure. Here in this inhospitable environment, a myriad of organisms communicate and navigate through a wide variety of beautiful and ephemeral displays of coloured lights.

Some interesting examples include the 246 species of lanternfish (members of the Myctophidae family), which live deep in the all of the world's oceans, making up an astonishing 600 million metric tonnes of biomass. This is about 10 times the world's yearly catch of fish! These abundant but small (most are just 6 cm long) deep-sea fish play an essential ecological role as food for larger fish. Myctophum punctatum, pictured below, lives a vertically mobile life. It rises over a kilometre to reach waters near the ocean surface every sundown, following the also vertical daily migration of zooplankton, its food source.

Lanternfish (Myctophum punctatum)

Although not readily visible in the photograph above, most of these species luminesce through photophores (light-producing organs) arranged in rows along the belly (they are not the reflective blue upper dorsal spots you see above). Photos of its bioluminescence are rare but the digital model of this species shown below offers an idea of how these bioluminescent photophores are arranged.


3D Digital  Model of Myctophum punctatum

Myctophum punctatum uses its bioluminescence as a special type of camouflage called counter-illumination. The fish regulates the brightness of the bluish light emitted by its photophores to match the blue wavelengths of the faint sunlight light streaming from above. This masks its silhouette from predators swimming underneath it. Some species also emit green or yellow light, which might be used for communication or courtship.

Just as fascinating are the anglerfish, an order of fish (Lophoformis) comprised of more than 200 species. All anglerfish are carnivores that use bioluminescence as a type of fishing lure. Representatives of these macabre-looking fish are shown below.

Masaki Maya et al., Wikipedia

Representatives of Anglerfish Order Lophoformis

Most, but not all, of these species live deep in the ocean, where the water is pitch black, extremely cold and under intense pressure. In these species, a piece of dorsal spine has evolved into a protruding "fishing pole" tipped with a luminous bulb. Not all individuals of each species have a fishing pole, however. Those that possess one are all females. Males are much smaller and have no need to fish for food because they are completely parasitic upon the females. (Click this link: what you first see is awesome!) Bulb-less, a male latches onto a female with its sharp teeth and eventually fuses with the female's body. It connects to her bloodstream, and eventually loses its eyes and internal organs, everything except its testes. Females are known to carry more than six males on their bodies at one time.

Some rare footage of an anglerfish mating pair reveals that these strange fish not only have bioluminescent lures but a surrounding net of fine bioluminescent filaments as well. They were filmed in 2016 at 792 metres deep off Portugal's coast.

The 3-minute video above was published by Science Magazine in 2018.

A vast number of species (in the thousands) that use bioluminescence, as well as a large number of variations in the chemical reaction that produces the light, suggest that bioluminescence evolved independently many different times throughout history. Experts estimate that it evolved at least 40 different times starting at least 150 million years ago, near the beginning of the Cretaceous Period.

In fish alone, bioluminescence evolved at least 27 different times. All this convergent evolution is a testament to its usefulness. Bioluminescence is a very useful jack-of-all-trades system, having numerous functions such as lures for prey, predator warning systems, mate attraction and communication. Bioluminescence is autogenic in many fish, which means that the animal itself produces the light. Examples are the lanternfish, which emit light from light-production organs called photophores. Many other marine species have instead left the job of light production to internalized luminescent bacteria. To further enhance this symbiotic relationship, they have co-evolved mechanisms within their bodies to turn these bacteria off and on at their will. They evolved a handy light switch for them in other words. Anglerfish belong to this group. Their mutually beneficial relationship with their luminescent bacteria is especially interesting. Recent genetic studies reveal that these bacteria have lost almost half of their genome compared to their free-living close relatives, an example of adaptive gene loss, a use-it-or-lose-it principle in genetic evolution. The bioluminescent bacteria appear to be able to swim in and out of the bulb freely using their flagella, but they have lost all the genes associated with sensing and digesting food sources. The fish, instead, supply all the cellular nutrients the bacteria requires while the bacteria provide light to help the fish lure food.

Perhaps one of the eeriest examples of bioluminescence is foxfire or faerie fire. Produced by some species of fungi in decaying wood, the light might be used to attract insects to spread the fungus? spores or as a warning to any foraging animals nearby. Omphalotus olearius, by day, is an ordinary looking (but poisonous!) European mushroom that looks much like a chanterelle:


Fungus Omphalotus olearius During the Daytime

By night, it reveals why it is also called the jack-o-lantern mushroom. Its orange gills glow an eerie green. It's easy to imagine faeries holding their mysterious night-meetings here.

Noah Siegel;Wikipedia

Omphalotus olearius at Night

How Does Bioluminescence Work? Where Did It Originally Come From?

Like a glow stick, bioluminescent animals utilize a chemical reaction that produces light. Because bioluminescence evolved independently in a wide range species across the evolutionary spectrum, different species use different chemical reactions to make light. The general mechanism, however, is the same throughout. It involves three essential ingredients: a light-emitting molecule, an enzyme, and molecular oxygen, O2. Inside the light-emitting cell, a special protein reacts with a charged oxygen ion and undergoes an oxidation reaction. During the reaction, an intermediate molecule in an excited state is produced. What does this mean? A molecule enters an excited state when one or more electrons in its atoms absorb enough energy to jump outward to a higher energy quantum state. Light is produced during the reaction when an excited electron in the intermediate molecule emits a photon of visible light. By dong so, it loses just the right amount of energy to drop to its ground (resting or lowest energy) state. This scenario isn't unusual in biochemistry. These chemical reactions often involve a short-lived intermediate molecule. It is usually highly reactive and it is often in an excited state. What is unique here is that the intermediate complex releases its excitation energy is as a photon rather than sequestering it in the potential energy of one or more new chemical bonds. The special protein in this reaction is generically called a luciferin (which means a light-emitting substance*). The light-emission reaction is catalyzed by a protein enzyme, generically called a luciferase. The trigger for the reaction can be a mechanical, neurological or chemical change.

*I was so curious I had to look up "Lucifer." How is this word also the word for devil? According to Wikipedia, Lucifer  originated as the Latin word for "bringer of light?" and as a metaphor for "morning star," which, in Christianity, also appears to have been the original name for Satan before his fall from grace. It then became a common byword for Satan or Devil.

Luciferin in Many Marine Organisms

The shape and size of the luciferin molecule varies widely among the various bioluminescent phyla. However, some luciferins are more commonly found than others. For example, a luciferin called coelenterazine is found in most bioluminescent marine organisms. It may have evolved into its luciferin role from an earlier role in detoxifying cells. Oxygen is essential for most life on Earth, but inside cells it is a hazardous molecule, especially when it is in the form of an even more reactive oxygen free radical. Coelenterazine is a strong anti-oxidant. It reacts with free oxygen radicals and neutralizes them before they can react with and seriously damage proteins and DNA inside cells. This might have been coelenterazine's original function. When ancient ancestors of marine organisms moved down into deeper sea habitats, where it is cold and oxygen levels are low, the metabolic rate of organisms would have gone down, requiring less intercellular coelenterazine. It was a molecule, already present, ready to evolve into its new luciferin function.

The enzyme luciferase, which catalyzes the oxidation of the luciferin, also varies widely. Each unique luciferin/luciferase system represents a unique evolutionary origin of bioluminescence. How did these molecules arise inside living cells? As suggested above, they appear to have mutated over time from biological molecules already present in the organism. Molecules very similar to the luciferin molecules seem to have already been present in many non-luminescent living organisms, doing some other function.

Luciferin in Some Dinoflagellates

Another intriguing example to illustrate this theory is a bioluminescent marine dinoflagellate. The luciferin used by most of these tiny unicellular organisms closely resembles chlorophyll, an even more ancient and biologically important molecule. Chlorophyll is also present because many dinoflagellates are photosynthetic. In some of these species, luciferin even seems to retain some of the light-absorption function of chlorophyll. Some species, such as P. lunula, bioluminesce only at night, and after sunny days they glow more brightly. The luciferin of this species might be a photo-oxidized chlorophyll. This chlorophyll molecule fluoresces blue when it is exposed to UV light (a part of sunlight) but the fluorescence eventually stops when all the chlorophyll molecules become degraded by photo-oxidation (oxidation in the presence of sunlight). Once it is photo-oxidized, it?s function changes and it becomes a bioluminescent luciferin, which glows blue. Even though this species offers a tantalizing clue about the origin of bioluminescence, researchers know that several other bioluminescent dinoflagellate species don't use this biochemical mechanism.

Luciferase in Fireflies

Luciferase, like luciferin, seems to have evolved from molecules that were already present in the organism. Over time, the molecule's function shifted toward to light-production, offering a new selective advantage to those individuals. This process of evolution toward bioluminescence is still far from understood but some headway has been made from studying coelenterazine luciferin evolution in marine organisms and by studying the chlorophyll-like mechanism for bioluminescence in a dinoflagellate species. Luciferases in beetles such as the firefly have also been the subject of close study, and again, the bioluminescent machinery might be borrowed from the cell's general biochemistry machinery. The firefly luciferase enzyme seems to have evolved from another enzyme called AMP-CoA-ligase. This means that somehow this enzyme evolved from breaking down fatty acids into an oxygenase/light-production function. Just to note here, the original ligase with its original function still operates in the organism. Its ligase role is still essential to its cellular biochemistry, especially in a number of regulative cellular functions. We can assume that the change in function toward producing light gave bioluminescent beetles such as fireflies a significant evolutionary advantage over their non-bioluminescent relatives, by enhancing their reproductive success. The numerous evolutionary paths of luciferin and luciferase are wonderful examples of Mother Nature re-purposing items at hand into new and amazing living tools.

Two things are fairly certain: First, bioluminescence evolved independently at least 40 separate times and second, it first evolved many millennia ago. There is evidence that it evolved in marine fish between 150 and 60 million years ago. In just under half of these marine fish, bioluminescence evolved not inside of them but in their symbiotic bioluminescent bacteria.

Evolution of Bioluminescence in Bacteria

It's possible, but not proven by any means, that bioluminescence first appeared in bacteria living in the early Cretaceous period (about 145 million years ago). Aerobic bacteria evolved once Earth accumulated enough oxygen in the atmosphere to support them. These bacteria had a metabolic advantage over their anaerobic cousins. They could now use the high-energy chemical bonds in oxygen to oxidize glucose (cellular food) during cellular respiration. Those products could then be used to make ATP (adenosine triphosphate), a very important energy molecule used by all cells.

Aerobic bacteria, and most, but not all, living organisms, also use oxygen to obtain energy from other molecules in addition to glucose, such as from fatty acids. Fatty acids, which I very briefly mentioned earlier, are all-around energy storage molecules in aerobic cells. Their energy, too, can be ultimately captured in ATP. When the aerobic pathway isn't used in the cell, anaerobic processes use fatty acids to make a variety of important molecules like phospholipids, messengers and hormones. Even the aerobic eukaryotic cells of multicellular animals carry out anaerobic processes.

Early aerobic bacterial species, living about a billion years ago at the end of the Great Oxygenation Event, likely lived in a changing environment where oxygen levels went up and down. This environmental stress might have been the first trigger for the evolution for bioluminescence. When oxygen levels fell, a mutation in an enzyme called riboflavin oxygenase (think of this molecule as an early luciferase) might have allowed these organisms to oxidize aldehyde molecules, which would accumulate under those conditions. By oxidizing the aldehydes, they could be turned into useful fatty acids. A supply of these molecules under low oxygen conditions would be a great evolutionary advantage. These bacteria would have that extra energy boost to keep reproducing and accumulating in numbers until oxygen became plentiful again. If some of these mutants also used another molecule, the already abundantly present reduced flavin mononucleotide (FMNH2) as a substrate or cofactor (an early luciferin in other words), then they could have been "accidentally" luminous, because this reaction creates an excited intermediate molecule that gives off light when it returns to its ground state.

If these early individuals were also light-sensitive, as many bacteria are, luminance among them might have been very useful in helping individuals recognize the presence of others nearby. It might have offered them a brand new selective advantage by helping them to disperse and/or colonize more easily. In bacteria, this behaviour, called quorum sensing, allows bacteria to detect and respond to their neighbours through gene regulation. The bacteria can switch genes on or off in order to optimize the population for changing conditions. Offspring might start producing a biolfilm, for example, so they can stick to an optimal rock surface. Or the population might switch to spore encapsulation when conditions become harsh. You could think of this new advantage like a military troop receiving radio communication on the battlefield - very useful under quickly changing conditions. All bioluminescent bacteria have a few characteristics in common: they are rod-shaped, gram-negative, they have flagella to move around with, and most importantly to the evolution argument I just laid out, they are all facultative anaerobes, which means they can live and grow when oxygen levels are high and also when they are low or zero.

The Genetics of Bioluminescence

Even though a single general pathway toward bioluminescence might have originated in facultative anaerobic bacteria, the chemistry of bioluminescence in bacteria varies depending on the bacterial strain or species. This suggests that bacterial bioluminescence evolved independently numerous times just as it did in other phyla. While the chemistry each time might be unique, all the luminescent bacterial species share the same gene sequence called the lux gene sequence, or lux operon, again offering an argument for the evolution of a single general pathway toward bioluminescence, after which the chemistry evolved independently numerous times. The lux operon might be the code left over for an ancient DNA repair system. Now this sequence codes for all the proteins involved in the luminescent mechanism. It is a short fragment of DNA just 9 kilobases long, and it contains just 5 genes that code for the proteins required for bioluminescence. A few additional genes regulate the operon (they turn it on or off). Several other enzymes, substrates and co-factors used in the production of light are already present in the cell. This short gene sequence can be isolated and inserted into normally non-luminescent bacterial and eukaryotic cells to make them bioluminescent. A non-luminous bacterium, such as Escherichia coli for example, can be transformed into a bioluminescent one simply by the insertion of the lux gene sequence. As you can imagine, this has a myriad of potential uses in medicine such as imaging and in research.

Bacteria tend to evolve fast because their life cycle is short and there are many of them. If a random mutation in the genetic code introduces a new advantage, survivors pass it on and it quickly increases in the population. The lux operon sequence is strongly conserved (which means it stays much the same with few surviving mutations) among bioluminescent bacteria. This suggests that it must have a significant selective advantage even though it is an expensive option for a cell to choose. There is a very high energy cost to emit light. A green light photon, for example, has about the same energy as the chemical bonds of 8 ATP molecules. That is a significant energy commitment for a microscopically small organism like a bacterium.

Bioluminescence Chemistry

Let's focus now on the light-producing reaction itself. The light-emitting reaction in bacteria has been studied extensively. Reduced riboflavin phosphate (FMNH2) and a long-chain fatty aldehyde (RCHO) are both oxidized, and oxygen diffused from the environment into the cell is the perfect oxidizer to do the job. During this oxidation reaction, blue-green light is emitted. The reaction is catalyzed by various enzymes called luciferases. The luciferase used depends on the species. FMNH2 and RCHO are already present in all aerobic bacteria (and eukaryotic cells) because they are part of the electron transport chain.

I was inspired to research bioluminescence after exploring the electron transport chain in the previous article, Ozone. This chain is part of the process of aerobic cellular respiration. "Aerobic" means requiring oxygen. Through aerobic cellular respiration, cells use food and oxygen and turn it into the energy needed to grow, multiply and move. The electron chain carries out one of the processes of aerobic respiration, called oxidative phosphorylation. Its main purpose is to make an important energy molecule called ATP.

The electron transport chain in bacteria (or in mitochondria in eukaryotic cells like ours) transfers electrons from donor molecules to acceptor molecules in a series of redox reactions. It is coupled with the movement of protons (H+ ions) pumped back across the cell membrane so that the cell and its environment remain neutrally charged. The chain drives the production of energy-rich ATP and it oxidizes a variety of enzymes and other proteins along the way. The final electron acceptor is oxygen, which is the perfect molecule for this because it is a powerful oxidizer (or "electron-grabber" if you want).

In non-luminescent bacteria, FMNH2 (riboflavin phosphate) simply diffuses into the cytoplasm. In bioluminescent bacteria this is where the luciferase enzyme and enzymes that catalyze the creation of the intermediate complex are also present. They channel FMNH2 into forming part of a light-emitting complex. In the process, FMNH2 is reduced to FMN and a long-chain fatty aldehyde (RCHO) is reduced to a carboxyl acid (RCOOH). ("R" is organic chemistry shorthand for any carbon-hydrogen group)

This is what the generic reaction looks like:

FMNH2 + RCHO + O2 → FMN + H2O + RCOOH + hv (490 nm)

The products of the reaction are water, RCOOH, a carboxyl acid, and flavin mononucleotide (FMN), which is an electron carrier in the electron transport chain of every living cell. Energy is released in the reaction in the form of a blue-green photon [hv (490 nm)].

During the reaction, the about-to-be-reduced FMNH2 binds to the luciferase enzyme. An enzyme is a biological catalyst. It increases the rate of the reaction but it isn't consumed during it. The luciferase catalyzes the reaction by reacting with oxygen and then interacting with the aldehyde to form a fairly stable intermediate complex in an excited state. It decays, or returns to ground state, slowly. This means that light can be emitted over a significant period of time by numerous complexes rather than just in a single brief flash (although some species do flash). The particular enzyme utilized by each specific species can have a significant effect on the decay rate (duration of light production) and the turnover rate (how soon it can glow again) of this light-emitting complex.

Although bioluminescence might have evolved during exposure to low-oxygen environments, molecular oxygen (O2) is essential for the light emission reaction pathway. Marine organisms get it from the oxygen in seawater. Land organisms get it from air. The reduction of oxygen ultimately transforms potential chemical bond energy into light energy.

Although this reaction produces blue-green light, bioluminescence can show up in a variety of colours. Blue-green or blue light is most common. In a marine environment where light levels are very low, blue light is the only visible wavelength short enough to remain streaming through the water after longer wavelengths have been scattered by water molecules. This makes light in the blue spectrum very useful for communication in dark deep marine waters.

Single mutations in luciferase can modify the chemical binding sites on the light-emitting complex just enough to distort the emission colour. Some bacteria also carry fluorescent proteins that change the emission colour, for example, to yellow. These proteins absorb blue light and re-emit it as less energetic wavelengths like yellow. A few organisms emit red. In each case, the basic chemical reaction that produces the light remains the same.

Bioluminescence Is Useful To Humans

Now that the mechanics of luminescence in various species are being worked out, an increasing number of new technologies are being developed to take advantage of it. Bioluminescent imaging, for example, allows scientists to non-invasively study biological processes while they are taking place inside live subjects. One especially promising new idea is to use it to track the progress of cancer metastasis in the living body in order to understand the progress of cancer better. Bioluminescent DNA machinery, the lux operon, can be inserted into various types of non-luminous cells.

To start, the lux operon can be spliced into a virus's genome. The virus then infects a cell and inserts that DNA into the cellular DNA. The cell can then translate and transcribe the luciferase/luciferin proteins that emit detectable light. If, for example, the lux operon is inserted into the cancer cells of a primary tumour of a laboratory animal such as a mouse, that tumour will light up and glow. As the cancer cells spread into the blood stream and invade other organs over days and weeks, the process can be observed and studied simply by observing where the mouse glows. Of course, visible light can?t pass through an animal but researchers got around this problem by culturing bioluminescent cells and selecting for lux operon mutations that emit near-infrared light rather than blue or yellow light. Cancer cells emitting this wavelength of light can be imaged and tracked over time using an infrared camera placed outside the mouse's body. This kind of imaging can be so sensitive it can detect the glow from a single cell. It is a way to see exactly where cancers spread in the living body over time.

Bioluminescence can also be used as a very sensitive assay in genetic research. Researchers can take the lux genetic code for firefly luciferin/luciferase, for example. This genetic sequence can be isolated and cloned onto any DNA sequence of interest and then inserted into a virus, which has all the cellular machinery to make proteins from the inserted genetic code. The protein that is transcribed can be measured itself or its enzymatic activity can be measured by the intensity of its bioluminescence, in this case, a yellow glow. Each protein molecule transcribed from the DNA sequence of interest, perhaps it is code for a specific enzyme or a regulatory protein, now has a glowing tag attached to it. It is a very sensitive assay.

The idea of using tiny glowing markers to visualize specific protein molecules in a petri dish or in tissue or in a living organism is not new. I already briefly mentioned green fluorescent protein (GFP), a breakthrough assay protein developed in 2008, and now widely used in medical and scientific research. The genetic code from GFP is also used as a genetic assay, but in its case the marker protein fluoresces green when exposed to blue to ultraviolet light.  Both bioluminescent protein complexes and fluorescent proteins emit visible light as the result of an excited electron in an energized molecule returning to its ground state. The bioluminescent marker protein, however, luminesces in total darkness and does not need to be in a lit environment. This means there is no background to eliminate when measuring the light emitted from an assay result and this makes it up to a thousand times more sensitive than GFP assays. It means as well that extremely small changes in light emission now become measurable by using ultra-sensitive cameras to detect changes too small for our eyes to detect.

A Glimpse at Future Technologies

Bioluminescence has fascinated humans for millennia. Technologies that use bioluminescence are probably still in their infancy and the future looks bright (sorry). Think of trees gently lighting future streets at night  à la Avatar. What a way for a city to go green and save electricity costs! The myriad possibilities of bioluminescent technologies are limited only by the imagination.

Monday, October 8, 2018


Ozone is a duplicitous character. We hear that it is bad for us (as a component of air pollution) and we hear it is good for us (as an atmospheric layer that protects us from harmful solar UV radiation). Which one is it? Where does it come from? How does it work? In this article I hope I can dispel some ozone mystery. The journey will take us through complex and interesting terrain: atmospheric chemistry, biochemistry and free radical chemistry.

The word ozone comes from a Greek word meaning "to smell," so named because this clear pale blue gas has a peculiar odour. You may have smelled ozone while operating a poorly working (sparky) appliance or you might have noticed it as that unique fresh scent right after a thunderstorm. I find it a bit sharp on the nose and weirdly pleasant.

Ozone is produced on Earth in three ways. First, when air is subjected to an electric spark such as lightning, ozone is created.  Second, ozone is produced as a byproduct of fuel combustion - from car engines, forest fires and industry. If ozone is present in high enough concentration over time, it is harmful to animals, plants and humans. This is called ground level, or tropospheric ozone, after the atmospheric layer in which it exists. A third method of production is natural and it occurs at very high altitude. Ozone is created when ultraviolet (UV) light from the Sun passes through oxygen in the stratosphere. This ozone forms a protective blanket that strongly absorbs UV light, a harmful form of radiation, preventing it from reaching Earth's surface. To see where the troposphere and stratosphere exist, the layers of Earth's atmosphere are shown below. The troposphere is the bottom layer in which we live. This relatively thin layer contains about 80% of the atmosphere's mass and almost all of its water vapour. It's where all of our storms occur. The stratosphere is the almost cloud-free layer just above it. Air here is much thinner. Atmospheric pressure here is about 1/1000 that at sea level. Jets rarely fly above the lowest part of the stratosphere.

Ozone Is Composed of Oxygen Atoms

Ozone is an allotrope of the element, oxygen. Oxygen allotropes are different ways in which oxygen atoms bond together. Oxygen can exist in a highly reactive atomic form, O1, or as the familiar stable colourless O2 gas we breathe in. Liquefied oxygen gas is pale blue. Oxygen can also exist as unstable and reactive ozone, O3, a pale blue gas or dark blue liquid under pressure. Oxygen can even exist as a dark red metallic solid, O8, under immense pressure.

Oxygen is a reactive element, in any allotropic form. O2 binds readily with most other elements and compounds to form oxides - chemical compounds that contain oxygen atoms. About half of Earth's crust consists of oxides and about one fifth of our atmosphere consists of O2. Having so much of this reactive gas in our atmosphere indicates that our planet bears life. No abiotic (non-life) processes are known to constantly replenish oxygen that is constantly sequestered into oxides by rock and minerals. The source is green plants. Green plants release oxygen gas into the atmosphere as a waste product of photosynthesis, a process that uses the Sun's energy to grow and make food. Oxygen-breathing animals like us evolved to use oxygen. We can trace our origins to ancient unicellular life forms that gained the ability to utilize this chemically reactive gas in order to power a series of reactions called cellular respiration. This process turns the food we eat and the air that we breathe into the power to move, grow and repair our bodies. Plants in turn utilize our waste gas, carbon dioxide, along with the Sun's energy, to grow and make food for us, in what is a wonderfully elegant symbiotic partnership.

Why Is O2 So Important to Life?

All of the oxygen allotropes mentioned earlier are present on Earth naturally. Each allotrope has different physical properties and chemical reactivity due to the unique structures and strengths of the chemical bonds between the atoms. O2 is the most chemically stable oxygen allotrope at atmospheric pressure and temperature. Essential for life for animals, fungi, protists and some bacteria, O2 has a unique electron configuration that keeps it stable in the air but reactive enough for life to exploit.

In O2, two oxygen atoms are bound together by a covalent double bond. This means that two pairs of electrons are shared between the two atoms. The diagram below gives you an idea of how this works. Electrons are the small green and purple circles below. The large central circles represent atomic nuclei. The oxygen atom belongs to group 6 on the periodic table, which means it has 6 outer electrons, available for bonding. These are called valence electrons.

Each oxygen atom has 8 electrons, 6 of which are valence electrons. The valence electrons have the same energy so they belong to the same (outermost) shell below. All atoms are most stable when the valence energy shell is full, with 8 electrons. Energy shells are shown as rings in the simple O2 electron shell diagram below. By sharing two pairs of electrons, two oxygen atoms are stabilized in a molecular structure that fills each valence shell. By bonding, each atom can attain 8 valence electrons.

The diagram left is a very simple way of looking at the bond in an oxygen molecule. This diagram is helpful but it doesn't show us one important fact. It doesn't show us how the electrons pair up with each other, as electrons tend to do. The oxygen molecule is quite unusual in fact because it has two unpaired valence electrons. These unpaired electrons explain why O2 is so useful to life. We can show this additional fact by drawing a modified Lewis diagram:

"A" to the left represents an ordinary Lewis diagram of O2. There is a double bond between the two oxygen atoms, drawn using two parallel lines. The un-bonded electrons of each atom are shown as two electron pairs. It makes sense - we learn early on in chemistry that electrons like to pair up. But it isn't quite accurate. If we refine our viewing lens once again, this time from a Lewis structure to a more complex and accurate atomic orbital representation, we discover that the double bond is actually composed of two orbital bonds: a sigma bond and a pi bond. An atomic orbital is a three-dimensional shape outlining where a particular electron might be orbiting a nucleus. This updated version takes into account that electrons are actually quantum wave functions. This model describes electron energy in much better detail and this helps us understand the bonding behaviours of the atom.

Every single chemical bond consists of one sigma bond. It is the strongest bond and it is simply the head-on overlapping of two atomic orbitals. A double bond has an additional pi bond. This bond is the sideways or lateral overlapping of atomic orbitals and it makes the bond stronger. A triple bond, stronger still, consists of a sigma bond and two pi bonds. To help you visualize these orbital-overlapping bonds, check out the simple diagrams on this site.

O2 bonding is rather unique. It is a double bond but in this case, the pi bond acts like two half-pi bonds plus two unpaired electrons. Take a look again at right hand drawing in the diagram above. The unpaired electrons are shown in red for emphasis in "B". This unusual configuration leaves two unpaired electrons with equal energy. Having two unpaired electrons allows O2 to reach the lowest potential energy state possible. This is something all atoms and molecules tend to do. If some heat were applied to the O2 molecule, those two electrons would pair up and the molecule would have slightly more potential energy. What we learned in school still holds up: these two unpaired electrons "want" pairing. This means that an oxygen molecule greedily accepts electrons from other atoms and molecules. It is chemically reactive in other words.

Most molecules tend to have paired electron spins. O2's unpaired electrons don't match up well with the valence electron pairs of other molecules. They are an awkward "third wheel" in the interaction. The consequence of this is that atmospheric O2 reacts slowly with most other substances, rather than rapidly. This is good because otherwise the oxygen in our atmosphere would trigger spontaneous combustion. An example in nature is the gradual process of rusting, the oxidation of iron exposed to air, into ferric oxides (rust). Oxygen is an oxidizing agent, which means it causes other substances to lose electrons. By doing so, oxygen itself gains electrons. This makes sense when you look at its two unpaired electrons. They "want" electrons so they can become pairs. The word "oxidation" was coined by Antoine Lavoisier, while observing reactions with oxygen. It is a bit of a misnomer because these types of reactions, more accurately called redox (reduction/oxidation) reactions, simply involve electron transfer. They don't have to involve oxygen.

How do our bodies utilize oxygen's unusual unpaired electrons? Mitochondria, the tiny "power plants" inside our cells, use oxygen as a final and powerful electron acceptor along a string of reactions called an electron transport chain. An electron transport chain is a series of redox reactions. Electrons are transferred from one molecule to the next. Differences in Gibbs free energy (chemical energy available to do work) between the reactants and the products drives this process forward. The beauty of this set-up is a) it's spontaneous and b) it transfers chemical bond energy to a molecule that can store and readily release it when required. Along the way, a molecule called ATP is produced. ATP (adenosine triphosphate) is the all-important energy storage molecule for all life - plant and animal. Like a tiny battery or fuel cell, ATP powers almost all cellular reactions. Driven backwards, the electron transport chain "burns up" ATP to provide energy for growth and for mechanical energy such as a sperm's flagellum or a contracting heart muscle.

Trivalent Oxygen, Ozone, O3

Now that we understand the chemical nature of O2, how does O3 compare? Ozone is a bent molecule. It has a triangular shape like water, H2O. The three oxygen atoms bond with a double and a single bond that resonates back and forth. Ozone is an example of a resonance hybrid.

The O-O bonds are a hybrid between a single sigma bond and a double sigma-pi bond. This means that the bond strength is in between that of a double and a single bond. Ozone's hybrid bonds are slightly weaker than O2's double bond.

Unlike its somewhat tamer cousin, ozone is a chemically unstable molecule. A resonant structure tends to stabilize a molecule but it is not enough to make ozone stable. The valence electrons in ozone are shared across three nuclei rather than two. Six valence electrons each are fighting for space while packed into a bent shape. Electrons with opposite spins like to pair up but electrons in general don't like to be too close to one another. This bent shape is the most stable lowest-energy arrangement possible but it still has high potential energy and that means it is unstable.

The hybrid bond structure means that the valence electrons in ozone are delocalized. These delocalized electrons spread out to form a loose molecular orbital cloud. Their enhanced motility allows them to react more readily than the localized electrons in O2 do, even though O2 has reactive unpaired electrons and ozone does not.

Ozone is one of the strongest oxidizers known, much stronger than O2. Because ozone is unstable, it readily decomposes into stable O2 gas and extremely reactive chemically unstable lone oxygen atoms. These lone atoms are the key to why ozone is such a strong oxidizer. Although these atoms have 6 valence electrons each, they don't all pair off. Two form pairs and two exist as lone electrons. Those two lone electrons mean that these atoms "want" electrons very intensely in order to stabilize themselves. They will immediately react and borrow electrons from almost any other substance they come across. Ozone is a much stronger oxidizer than O2 because O3 is less chemically stable. O3 offers up lone oxygen atoms that form as soon as ozone decomposes. Lone oxygen atoms are much stronger oxidizers than O2 oxygen molecules, because they so unstable (and therefore reactive).

O3 is formed when O2 reacts with highly reactive atomic oxygen, O1 (or just O). Fee atomic oxygen reacts and disappears almost instantly from Earth's lower atmosphere, but in the stratosphere it is continuously replenished. Stratospheric O2 is bombarded by UV (ultraviolet) radiation from the Sun, cleaving its bond into two free O1 atoms. In the stratosphere, O3, O2 and O1 all exist and are all part of a cycle. In low-Earth orbit, far above the stratosphere, the very sparse atmosphere is almost entirely composed of atomic oxygen, O1. This diffuse but highly reactive gas corrodes all the outer materials on spacecraft that pass through low Earth orbit. It is a significant challenge that all space agencies must take into account.

In the stratosphere, the O2 + O → O3 synthesis reaction is triggered whenever O is available. O is created when solar UV radiation breaks apart the O2 molecular bond into two free oxygen atoms. Whenever atmospheric O2 comes into contact with free atomic oxygen it quickly combines into O3, ozone. Because ozone is a much more powerful oxidant (electron acceptor) than O2 is, it is much too reactive to be useful in any cellular electron transport chain. In fact, its oxidizing action makes ozone pollution a serious health hazard. It can damage respiratory systems in animals and cause tissue damage in plants.

Although O2 is less reactive than ozone, it too is an oxidation hazard inside living cells, and this is something life has learned to live with. Various intracellular sequestering processes reduce this hazard. It is an evolutionary trade-off between cell damage and oxygen's electron-acceptor powers.

Free Radicals

You may have heard of how bad free radicals are for our health. These mysterious-sounding chemicals are simply atoms or molecules (or even ions) that have an unpaired valence electron. Radicals are an important part of biochemistry and atmospheric science. We already explored a radical when we looked at the molecular bonding of O2. The O2 molecule is a di-radical. It has two unpaired valence electrons. We also came across free monatomic oxygen, which is another di-radical, and a much more powerful oxidizing agent.

Like O2, all radicals are reactive, some more than others depending on their stability. Radicals are always oxidants because they accept electrons. Inside cells, free radicals can cause oxidative stress. Oxidative stress is basically a disturbance in the normal intracellular redox balance. The mitochondrial (and in plants, chloroplast) electron transport chain is an ingenious natural invention, but it's not perfect. A few electrons always "leak" out of the chain and react directly with O2 at the end. This reaction creates negatively charged O2, which is a free radical called superoxide. This is the Lewis diagram for it:

This highly reactive charged molecule causes oxidative stress inside cells. It reacts with biologically important molecules such as DNA and proteins. Like a bull in a china shop, it breaks DNA strands haphazardly so they cannot replicate and transcribe accurately and it denatures proteins, so they can no longer function as enzymes, hormones, antibodies and so on. This microscopic damage gradually builds up at the cellular level and in the body as a whole, an overall effect we observe as aging.

Stratospheric Ozone

Stratospheric ozone is the "good" ozone. It makes all surface life on Earth possible. In the stratosphere, ozone forms and breaks apart continuously. When atomic oxygen (O) reacts with molecular oxygen (O2), ozone (O3) forms:

O + O2 + M → O3+ M

A significant amount of energy is released during this reaction. It requires an additional body (M), such as a non-reacting molecule nearby that can carry that energy away. There are two reasons why energy must be released. First, the chemical bond energy of O2 (498 kJ/mol) is slightly higher than that of O3 (445 kJ/mol) so some energy must be released. Second, the free oxygen atom in the reaction is in an excited (high-energy) state, and that energy must be released as well.

Excitation can be explained using atomic orbital notation. An orbital, once again, is a three-dimensional cloud where an electron can be found. When atoms and molecules react with one another, their outermost electrons interact to form or break chemical bonds. In its ground (lowest energy) state, the electrons in an oxygen atom occupy the three lowest energy orbitals available: 1s22s22p4. The lowest energy 1s orbital can hold two electrons; it's full. The next higher energy 2s orbital can hold 2 electrons; it's full too, so these two orbital clouds hold 4 electrons in total. Next, the p orbital starts to fill up. In an oxygen atom, it holds 4 of 6 possible electrons. Oxygen's outermost orbitals are those with n=2 orbital energy. These are the 2s and the 2p orbitals. These n=2 orbitals equip oxygen atom with a total of 6 electrons available to react chemically. These are the valence electrons. In theory, electrons could occupy any of a very large number of possible orbitals in any atom but in a ground state (lowest energy) atom, electrons always minimize energy by occupying the innermost orbitals possible. In an excited state, one or more electrons move outward into higher energy orbitals. An excited oxygen atom is most simply denoted as 1s22s22p33s1 where one outermost (valence) electron has jumped up to a higher-energy 3s (n=3 energy) orbital. This electron configuration contains one unpaired valence electron, which makes it a radical too. Oxygen radicals are denoted as O(1D). I won't go into the reason for the "D" here, but if you want to know, this NASA page explains it well. Not all radicals are in an excited state. Recall that a lone ground-state oxygen atom is a radical too; in fact it's a di-radical. The whole story of oxygen radicals can get pretty confusing. The important thing to remember is that all radicals are highly reactive because the unpaired electron always "wants" to pair up with a valence electron in another atom or molecule. When it does so, the system releases potential energy and stabilizes.

Atomic oxygen radicals (O(1D)) are very reactive because they are radicals and they are very energetic because they are excited. The unpaired electron in the valence shell of this atom combines rapidly with any O2 molecule it slams into to form ozone. In order to exist at any concentration in the stratosphere, energetic free oxygen atoms must be continuously made. They come from the photodissociation of molecular oxygen, O2. Photodissociation means the splitting of molecules by electromagnetic radiation, or light ("photo"). High-energy, and therefore short wavelength, UV radiation pierces the stratosphere and cleaves O2. UV photons with wavelengths shorter than 242 nm (nanometres) have enough energy to break the bond between two oxygen atoms in an O2 molecule. This energy corresponds to 498 kJ/mol. That's the bond energy of O2. The two oxygen atoms released in the reaction absorb some UV energy, leaving them in an excited state.

In the stratosphere, ozone cycles continuously, forming and decomposing O3. Both processes absorb harmful solar UV radiation, particularly of wavelengths shorter than 242 nm. This is why the ozone layer is a protective blanket against UV radiation. As we just learned, ozone absorbs short (<242 nm) wavelength UV radiation when it is produced. Ozone isn't chemically stable, so it doesn't stay around for long in the stratosphere. When it itself is bombarded with stratospheric UV radiation, it readily photodissociates back into O2 and O. The molecular bond energy of ozone is 445 KJ/mol, which is less than that of O2 (498 kJ/mol). This means that less energetic UV wavelengths will break ozone apart, those between 240 and 320 nm. The ozone photodissociation reaction formula looks like this:

O3 + UV (240nm -320 nm) → O2 + O(1D)

Excited free oxygen atoms from this reaction continue the cycle, creating ozone once again.

Stratospheric Ozone Absorbs Deadly UV Radiation

All of these reactions are fast; a whole cycle takes place in just over a minute. It is a very effective life-protecting "UV absorption machine" that converts UV radiation into thermal energy. That energy is carried by the excited fast-moving free oxygen atoms. The stratospheric layer, above the thermosphere, ranges from about 20 km in altitude in the tropics to just 7 km in altitude at the poles. It is a generally stable layer of air that ranges from about -51°C at the top of the troposphere to just -3°C at the top of the stratosphere. You expect the temperature to go down as you move upward through the atmosphere, but the thermal energy created by the ozone cycle is most active at the top of the stratosphere where incoming solar UV radiation bombards oxygen.

The Sun bombards Earth with all wavelengths of UV radiation (and other EM radiation as well). UV radiation ranges from 10 nm to 400 nm. Wavelengths shorter than 121 nm ionize air so strongly that they are absorbed long before they can harm life on the surface. An atom or molecule is ionized when it gains or loses electrons to form charged ions. Another short mini-lesson here: What makes an atom an ion is when the number of electrons doesn't match the number of protons, so the atom therefore has an unbalanced charge. High-energy UV photons have enough energy to cleave various atmospheric molecules apart into ions while the photons are absorbed in the process. A radical is an atom that has at least one unpaired electron. In this case the electron number may still match the proton number and in that case it isn't electrically charged, but it is very reactive. The charged superoxide radical we encountered earlier is both an ion and a radical.

UV radiation between 100 and 280 nm is deadly to almost all life on Earth. This is the wavelength range, especially 230 to 270 nm, utilized in special mercury, LED and xenon germicidal lamps. It kills almost all known microorganisms. Most microorganisms have not evolved protection against concentrated mid-range UV bombardment. It is the right energy to break apart chemical bonds in DNA, proving deadly. Rare exceptions are extremophiles and ancient bacteria that lived before Earth had a protective ozone blanket. These organisms oxidized iron and built protective "rust blankets" around themselves to shield them from UV radiation. There is evidence that photosynthesis evolved in these bacteria.

Fortunately for us, the most DNA-damaging UV range (130 nm and 260 nm) is completely absorbed by stratospheric ozone. However, a small amount of slightly longer wavelength UV radiation, between about 260 and 300 nm, does make it to the surface. This is the UV radiation (especially between 265 and 275 nm) that causes sunburns and can lead to deadly melanoma. It also causes eye cataracts and other eye damage.

As you might have noticed, it is within the range that is germicidal. How does it kill germs but not us? We and other multicellular life survive because, first of all, the natural solar bombardment of this UV radiation is far less intense than a concentrated beam from a lamp. Our cells therefore have a chance to repair the damage as it happens. Secondly, we have evolved some protection through our skin. A pigment called melanin absorbs UV radiation, directing it away from vulnerable cellular proteins and DNA. Our skin even makes use of some UV exposure (280 to 315 nm) to make Vitamin D.

Tropospheric Ozone

Tropospheric or surface ozone is the "bad" ozone. It is a respiratory irritant and it can cause plant tissue damage as well.

Ozone As Ground Level Pollution

Ground level ozone is a pollutant and it is a key ingredient in smog. A pollutant is a substance that is introduced into an environment that has undesirable effects on it or on the life that depends on it. We tend to think of pollutants as man-made but not all of them are. Some are created naturally such as volcanic dust and volcanic gases. Ozone is technically called a secondary pollutant because it is created in the atmosphere when - react in sunlight. These primary pollutants come from combustion engines in vehicles, from industry and from forest fires.

The reactions that create ozone occur best on hot sunny summer days, when there is plenty of solar (UV) radiation. High temperatures promote ozone accumulation by increasing the rates of reactions that form ozone and by reducing the ability of plants nearby to absorb ozone out of the atmosphere. Plants absorb a variety of air pollutants including as much as 20% of atmospheric ozone production. However, during heat waves, stressed plants close their stomata ( (epidermal pores) in order to conserve water and this means that they cannot absorb ozone and other pollutants.

Both NOx and VOCs come from natural sources as well as man-made sources. A significant amount of VOCs is released from coniferous forests, volcanoes and wildfires. NOx compounds are released during lightning storms and wildfires. There are many man-made sources of these pollutants, ranging from motor vehicle exhaust, oil refining, paints, insecticides and industrial solvents to chemical manufacturing, but most man-made NOx and VOCs come from motor vehicle exhaust. Motor vehicles are responsible for at least half of the concentration of these pollutants, especially in large cities even though catalytic converters have been mandatory since 1975 (at least in North America).

In Canada, ground level ozone advisories are issued when average levels per hour exceed 82 parts per billion. Toronto, for example, typically experiences about 10 ozone advisory days each summer. To see live readings for various Ontario cities, check this government website. In Edmonton, close to where I live, ozone pollution risk is usually low. Our cities are a bit smaller than Toronto but just as importantly we don't tend to experience summer days as hot as Toronto does.

This past summer saw Edmonton and the surrounding area blanketed in thick haze blown in from several large forest fires to the west in British Columbia. We experienced many consecutive days in August where air quality health indexes (AQHI) sat over 10+ (very high risk). The air quality health index measures the combined health risk of all fine airborne particulate matter as well as ozone and nitrogen dioxide. If at any time fellow Albertans want more specific information than an AQHI reading, check out this Alberta website that shows current Edmonton/central Alberta levels of ozone, NO2, fine particulate matter, sulphur dioxide and carbon monoxide. To get an idea of how major cities in Canada stack up internationally, this Canada government website compares the average annual ozone levels (in ppb) of various Canadian cities with selected international cities. Across the globe, tropospheric ozone levels range between less than 10 ppb over remote tropical oceans and over 100 ppb downwind of large metropolitan cities in hot weather.

Some ozone can enter the troposphere from the stratosphere through disturbances such as hurricanes that can draw some lower level stratospheric air downward. However, the vast majority of ground level ozone is created when it forms from reactions of precursor compounds such as VOC's and nitrogen oxides. Ozone is highly reactive so it leaves the troposphere quickly, but plants, animals and people downwind from large cities on hot days or downwind from large forest fires can face a significant ozone hazard.

In the stratosphere, we now know that ozone forms from the photodissociation of O2 into oxygen atoms, which recombine with O2 to form ozone. However, this reaction doesn't happen where plants, animals and humans live because short wavelength UV light (<242 nm) doesn't penetrate down into the lower troposphere. As in the stratosphere, the production of ozone requires atomic oxygen. Here, surface nitrogen dioxide (NO2) does the job of supplying it. Its photodissociation requires much less UV photon energy than molecular oxygen does - anything under about 420 nm (slightly more energetic than visible violet light) will work:

NO2 + UV (<420 nm) → O(1D) + NO

The oxygen radical produced in this case is not in an excited state. It will react with oxygen gas to create ozone:

O(1D) + O2 + M → O3 + M

However, in unpolluted air, there is no net production of O3 because O3 quickly reacts with the product NO to create O2 and NO2 once again, in a cyclic reaction (which isn't shown).

When other ozone precursors such as man-made pollutants like carbon monoxide (CO) and hydrocarbons such as methane (CH4) are also present in the air, net ozone build-up can occur. This is when ozone can spike to unhealthy levels. Some of these reaction mechanisms are extremely complex, but two fairly simple surface ozone production pathways, using carbon monoxide and methane as precursors, are fairly easy to show. Their formulae are written below.

Both reactions require hydroxyl radicals (*OH).

A hydroxyl radical, denoted *OH, is the electrically neutral form of the hydroxide ion (OH1). It has one unpaired electron as shown in the Lewis diagram below left.

Hydroxyl radicals are created when surface ozone is exposed to the longer UV radiation that reaches Earth's surface:

O3 + UV (240nm -320 nm) → O2 + O(1D)

The free radical oxygen atom produced then reacts with water vapour to create hydroxyl radicals and oxygen gas:

O(1D) + H20 → 2*OH + O2

Hydroxyl radicals are highly reactive and they are an important part of atmospheric chemistry. Denoted *OH, hydroxyl is sometimes called an atmospheric detergent because it reacts with many pollutants, decomposing them into smaller less harmful compounds. In this case, however, *OH is a step in producing a pollutant: ground level ozone. Compounds commonly present in combustion vehicle exhaust, such as nitrogen monoxide (NO), nitrogen dioxide (NO2) and hydroperoxyl radicals (HO2), serve as reactants and as catalysts. Catalysts in this case increase the reaction rate of surface ozone formation reactions. Faster production means that ozone can build up temporarily even though it is unstable.

Example 1: Carbon monoxide in an NO-rich environment:

CO + *OH → H + CO2
H + O2 → HO2
HO2 + NO → *OH + NO2
NO2 + UV (<420 nm) → O(1D)  + NO

Notice that the fourth reaction just above us is exactly the same reaction as in the natural O3 production reaction (the one that cycles and doesn't build up ozone). Here, however, the reaction takes place in polluted air where various pollutants catalyze ozone production. Ozone is therefore produced faster than it can be removed (which is by reacting with NO to create oxygen and nitrogen dioxide). The reaction scheme continues as the oxygen radical reacts with oxygen gas to create ozone:

O(1D) + O2 → O3

Example 2: Methane (CH4) in an NO-rich environment:

CH4 + *OH → CH3 + H2O
CH3 + O2 → CH3O2
CH3O2 + NO → CH3O + NO2
CH3O + O2 → CH2O + HO2
HO2 + NO → *OH + NO2
NO2 + UV (<420 nm) → O(1D) + NO
O(1D) + O2 → O3

In this case, both ozone and formaldehyde (CH2O) rapidly build up in the lower atmosphere. Ozone is gradually removed as it reacts with hydroperoxyl radicals. As ozone blows into non-polluted air where NO levels are low, it will react with HO2 generated in the reaction "line 4" above to create *OH radicals and oxygen gas. Further ozone depletion occurs when the *OH created then reacts with additional ozone to create new HO2 and more oxygen gas. Ozone is also deposited onto surfaces, where it can react with the surface it lands on. This is how plants are damaged by ozone. Ground-level ozone causes more plant damage than all other air pollutants combined.

Ozone pollution levels peak in the late afternoon, when solar UV radiation and therefore photochemical reactions peak. Ozone is much more likely to be a hazard near large cities and factories on long sunny days in calm air, rather than during short winter days, even though general air pollution levels might be similar or even higher as during a temperature inversion.

Ozone Is An Oxidation Threat To Our Bodies

Ozone is a toxic gas. That being said, we breathe in a tiny amount of it every day. In fresh unpolluted air at sea level, natural ozone makes up about 10-15 parts per billion (ppb) which means that every 15 billion air molecules will include an ozone molecule, on average. Our lungs breathe it in, handling it without noticeable damage. Highly polluted stagnant air, however, can contain more than 125 ppb ozone. Exposure to this much ozone over several hours or days can significantly harm humans and other animals. Long-term exposure essentially causes premature aging in our lungs. It can inflame lung tissues, cause throat irritation and shortness of breath, increase one's susceptibility to respiratory infections and it can aggravate asthma and COPD (chronic obstructive pulmonary disease). Ozone reacts with both the epithelial cells of the respiratory tract and with the molecules in the fluid that coats the tract, creating a variety of free radicals and other oxidant molecules that damage epithelial cells by causing oxidative stress. An enzyme released from the cytoplasm (cellular fluid) leaked from damaged epithelial cells attracts inflammatory cells, leading to reddening and swelling of the respiratory tract. This can in turn lead to difficulty breathing. Ozone also stimulates special nerve cells that exist in between the epithelial cells lining the respiratory tract. This stimulation causes the respiratory pathways to constrict. It also induces coughing and a reflex that reduces one's ability to inhale fully. All of these respiratory effects of ozone are reasons why it is a good idea to avoid strenuous activity outdoors during an air quality alert, even when you are healthy and especially when you already suffer from a respiratory problem. You can expect high ozone levels whenever pollution levels from combustion engines or industry are high or when you live downwind in the smoky hazy air blowing in from forest fires. Although individuals vary widely in their sensitivity to ozone, most of us recover completely from short-term exposure that lasts a few hours or less. Our respiratory tissues repair themselves quickly and they usually recover completely in about 48 hours.

While environmental ozone is a potent health threat, some cells in our bodies have actually evolved ways to use it and other similar oxidizing molecules to their benefit. For example, during an infection, activated white blood cells, called neutrophils, produce ozone and ozone-like oxidizing molecules. These potent oxidizers kill the bacteria invading our system by using a process that is sometimes called an oxidative burst. How it all works is still quite mysterious because, for one thing, it is difficult to examine what happens chemically during a process that occurs very rapidly inside living cells. Ozone and ozone-like radicals appear to be used in the creation of deadly nitric acid that is stored up inside tiny intercellular sacks called phagocytes. Phagocytes are like the trash compactor units of the cell. A phagocyte will engulf a bacterium into a nitric acid bath that denatures its DNA, killing it. It is also possible that the radicals themselves directly destroy the engulfed bacterial DNA.

Tropospheric Ozone Damages Plants

Plants species vary in their sensitivity to ozone but when ground level ozone exceeds 80 ppb for over four hours, plant damage can generally occur. This is the same ozone concentration that prompts human health advisories in Canada, so when we are at risk so are many of our plants. You will see broadleaf damage first show up as clusters of tiny reddish or purple dots in between the veins of leaves that are most directly exposed to sunlight.

Stippling on a red alder leaf caused by ozone pollution. Pat Temple, U.S. Forestry Service; Wikipedia
Often, leaves subjected to accumulating long-term exposure eventually turn autumn-like colours or brown prematurely and drop off. They basically succumb to oxidative stress. Plants look like they do at the end of their season, which makes it difficult to distinguish ozone damage that occurs in late summer. Plants already under stress, such as from drought for example, show more pronounced damage.

Here in Alberta, damage is first noticed on sensitive plants such as blackberries, ash trees and big-leaf lindens rather than more tolerant trees like spruce, pine and birch trees. Generally, the leaves on sensitive plants have more and/or larger stomata, pores that open to allow the plant to transpire (exchange gases) so they allow more ozone in. Ozone enters leaves like other gases do, through the numerous stomata. Once inside the leaf, ozone dissolves in the water inside the plant and reacts with other chemicals. It is a powerful oxidant that damages the photosynthetic apparatus inside the leaf. Once this damage happens, carbon dioxide levels begin to rise inside the leaf because it is not being consumed in photosynthesis. This stimulates the leaf to close its stomata, which further reduces photosynthesis. The plant, as a result, can longer make sugars effectively to maintain its health. There is evidence that leaves higher in antioxidants such as vitamin C have some resistance to ozone damage. By reacting with ascorbate (Vitamin C) in the watery cytoplasm inside the plant leaf cell, ozone is transformed into a variety of nontoxic products that the cell can handle.

Ozone damage to our global food supply is significant. A 2011 article suggests that global yield from three ozone-sensitive crops - wheat, soybean and maize - could be reduced by between 17% and 26% by 2030 based on projected upper and lower estimates of carbon-based emissions by the IPCC (Intergovernmental Panel on Climate Change). An effective way to reduce ozone crop loss is to move toward non-combustion green technologies in vehicles and in industry.

Ozone Threat is Close To Home

We might think of large international cities when we think of the threat of ozone pollution but the problem of ozone (and all air pollution) sits close to home. Alberta, famous for its clear blue skies, is also known for its oil, natural gas and coal production, and for its vast agricultural lands, all of which contribute to significant air pollution and, therefore, ozone pollution. In Alberta, we can expect higher ozone pollution downwind of our two major cities, Edmonton and Calgary, based on contributions of primary pollutants from vehicle exhaust and industries within city limits. However, ozone pollution can also be expected downwind of the oil sands in northern Alberta. The oil sands pump out between 45 and 84 tonnes of organic aerosols per day, a level comparable to that produced by the entire Toronto metropolitan area (about 67 tonnes per day). Organic aerosols are a poorly understood highly complex series of air pollutants, many of which interact with sunlight to create additional secondary pollutants, and which make up most of the fine particulate matter in air pollution.

Perhaps surprising is the fact that the smaller Alberta city of Red Deer has the worst air quality in Canada, according to numerous reports that came out in 2015. Current studies are still being done to figure out what the pollution consists of and where it comes from but the results so far seem to focus on two culprits - nitrogen dioxide and volatile organic compounds, compounds associated with industry and key ingredients of ozone production. Contributing to the problem is the fact that Red Deer sits in a bowl between river valleys, where air can sit and stagnate, and on hot summer days, one could expect high ozone levels as well.  Finally, if British Columbia continues to suffer from devastating wild fires every summer (the last two summers were record-breaking), Alberta will be blanketed by the smoke and haze as most wind flow is from west to east here. Fires tend to coincide with hot dry sunny weather, so Alberta will also suffer from seasonal ozone pollution.

Ozone is a fascinating Jekyll and Hyde type of molecule. Understanding how ozone works means understanding how chemical reactions work as well as how energy affects molecular interactions. The complex machinery inside our cells can make use of the redox chemistry that utilizes various oxygen allotropes but at the same time all living cells must protect themselves from the powerful oxidative activity of these same molecules. Ozone, originating from oxygen photosynthesized by plants, protects all life from deadly solar UV radiation. Yet, when it makes direct contact with the cells of life, it is a poison. An appreciation of the dual nature of ozone paves the way to an introduction of three challenging branches of chemistry: atmospheric and radical chemistry and biochemistry. It highlights how intimately these different branches are linked.