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.
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| Kelvinsong;Wikipedia |
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.
"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.
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:
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| DoSiDo;Wikipedia |
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 (https://en.wikipedia.org/wiki/Stoma) (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.
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| Attys;Wikipedia |
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.
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| Stippling on a red alder leaf caused by ozone pollution. Pat Temple, U.S. Forestry Service; Wikipedia |
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.
