Monthly Archives: December 2011

EPA’s New Ozone Rule: Part 4

This post is based on information provided by two websites:

  • Washington University in St. Louis, Missouri. Chemistry 152: How Does Ozone Form? Unfortunately, the link to this website has been broken and the website is no longer available. (A shame, as it was really well done.)
  • NASA Earth Observatory, Chemistry in the Sunlight: The Chemistry of Ozone Formation To view, click here

See also Geoffrey Tyndal, Peroxy Radicals: Big Players on Ozone Production, website hosted by the project Stratopheric Processes And Their Role in Climate (SPARC) (University of Toronto). To view, click here.

Special thanks to Dr. Tyndal who graciously answered my questions regarding ozone formation. He explained to me how the VOC molecule whose peroxy radical had surrendered a oxygen atom to nitric oxide could continue to form more ozone, allowing concentrations to grow.

See also the website Environmental Science Published for Everybody Round the Earth (ESPERE), Lower Atmosphere: Oxidation in the Atmosphere. To view, click here.

If I am misunderstanding or misinterpreting the information, readers are urged to correct me. First, I want to describe the chemicals that we will be discussing here:

Nitrogen oxides consist of some combination of nitrogen and oxygen. We will be discussing two here. The first is nitric oxide, whose molecule consists of one nitrogen atom bonded to one oxygen atom:


Nitrogen dioxide is a gas whose molecules consist of one nitrogen atom bonded to two oxygen atoms:


Volatile organic compounds (VOCs) are organic compounds (compounds containing carbon) that can exist as a gas or vapor in the atmosphere. Examples include natural gas, drinking alcohol, gasoline, fresh paint, nail polish remover, cooking aromas, women’s perfume.

Radicals are fragments of molecules that have one or more unpaired electrons. This makes them highly reactive: they are desperate to attach themselves to any molecule that will take them. They can have a positive, negative, or neutral charge. We will mention three types of radicals in this post, all of them neutral:

  • hydroxyl radicals consists of a hydrogen atom bonded to an oxygen atom, the chemical formula being (the mid dot [·] after the formula denotes electric neutrality):


  • Peroxyl radicals consists of two oxygen atoms. Normally, O2 is ordinary oxygen gas, but it can act as a radical, as we shall see.
  • Hydroperoxyl radicals consists of a hydrogen and two oxygen atoms bonded together:


  • In my last post, I wrote about the formation of stratospheric ozone that it is primarily produced by ultraviolet radiation from the sun. The ozone there does an excellent job in absorbing the radiation, so that very little ultraviolet radiation (of the type that forms ozone) reaches the ground. Ground-level ozone, by contrast, forms as the result of the chemical reactions of atmospheric nitrogen oxides. These nitrogen oxides are mostly formed by the burning of fuel in the presence of nitrogen and oxygen, the two most prevalent gases of Earth’s atmosphere. The high temperature of combustion fuses the nitrogen and oxygen into one molecule. This is why the exhaust of cars and coal-burning electric plants can be laden with nitrogen oxides. Lightning and chemical processes in the soil are also sources.

    There are many forms of nitrogen oxides; the two that are relevant here are nitric oxide (NO) and nitrogen dioxide (NO2). Only nitrogen dioxide directly produces ozone. A nitrogen dioxide molecule struck by a photon of sunlight (represented here as hv) will break down into nitric oxide and a free oxygen atom:

    NO2 + hv → NO + O    (1)

    The free oxygen atom is highly reactive and will quickly combine with a nearby oxygen molecule to form ozone.

    O2 + O → O3    (2)

    If that was all there was to it, the ozone problem would be much more manageable. Ozone is not a stable molecule: it reacts with surrounding chemicals to break down into ordinary oxygen. For example, ozone recombines with nitric oxide to reform nitrogen dioxide and ordinary oxygen:

    O3 + NO → NO2 + O2    (3)

    The amount of ozone would be much lower than it actually is, because it would be destroyed soon after being formed. There is a complicating factor, however.

    Volatile organic compounds (VOCs) can convert nitric oxide back into nitrogen dioxide with the help of the hydroxyl radical, a substance that composes a tiny proportion of the atmosphere (parts per trillion!) but can have a significant impact nonetheless. Suppose we represent a VOC molecule with the single chemical symbol R. Now VOC molecules usually have several hydrogen atoms. What we will do is exclude a single hydrogen atom from what is represented by R, such that the entire molecule is represented by the symbol:


    where H represents the single hydrogen atom.

    When the VOC molecule contacts a molecule of ordinary oxygen and a hydroxyl radical (OH·), the single hydrogen atom is replaced with a peroxy radical (O2·). A water molecule is also formed. The reaction is:

    R-H + HO· + O2 → RO2· + H2O    (4)

    When the VOC molecule with the peroxy radical attached encounters a molecule of nitric oxide (NO), the peroxy radical gives up one atom of oxygen to the nitric oxide, converting it into nitrogen dioxide:

    RO2· + NO → RO· + NO2    (5)

    The nitrogen dioxide is then free to form more ozone.

    But it doesn’t stop there. If the VOC molecule with the extra oxygen atom has another hydrogen atom (and it usually does), it can react with an oxygen molecule to form a carbonyl compound (a compound where an oxygen and a carbon atom share a double bond) plus a hydroperoxyl radical (HO2·). I represent the VOC molecule in the following equation as RCHO· to emphasize the presence of a carbon atom, a hydrogen atom, and an oxygen atom radical, and I represent the carbonyl compound as RC=O:

    RCHO· + O2 → HO2· + RC=O    (6)

    The hydroperoxyl radical can then react with nitric oxide to form the hydroxyl radical plus nitrogen dioxide, both of which can participate in further reactions to form more ozone:

    HO2· + NO → HO· + NO2    (7)

    This description is an oversimplification. There are other chemical pathways for ozone to form. Also, small amounts of ozone can drift down from the stratosphere. But this gives a good idea how much of ground-level ozone is created. The main point is that ozone creation is a cycle. Nitrogen dioxide reacts with ordinary oxygen to form nitric oxide and ozone. Nitric oxide is converted back to nitrogen dioxide by the peroxy radical in VOC molecules. It is then free to convert more oxygen to ozone, and on the cycle goes. The presence of VOCs in the atmosphere, therefore, maintains a level of ozone in that vicinity.

    Why is this information important to our question of ozone regulation? We need to sit down with industry and discuss what needs to be done to reduce ozone concentrations at ground level. Industry emits very little ozone; it does emit ozone precursors and these need to be controlled. By knowing how much these precursors need to be reduced, we can work with industry to set up best practices while minimizing the burden, trying to keep both necessary capital expenditures and additional operational costs as low as possible.

    Finally, it appears that a small amount of ground-level ozone (below 15 ppb, perhaps?) is actually beneficial, because it generates the hydroxyl radical. The hydroxyl radical has been called the detergent of the atmosphere, because it is very good at eliminating many atmospheric pollutants. The hydroxyl radical constantly needs to be replenished, because its atmospheric lifespan is only a few seconds. See the article written by Phillip Ball entitled Fast-Acting Atmospheric Detergent (starting with the fifth paragraph) in the journal Nature, November 3, 2000, which you can view by clicking here.

EPA’s New Ozone Rule: Part 3

To understand atmospheric ozone, it is important to know how it is formed. In a previous post, I pointed out how ground-level ozone affects humans differently than the higher-up stratospheric ozone (thus the quip, “Good on high, bad nearby”1.) The way they are formed is different as well. Stratospheric ozone is formed when solar ultraviolet radiation strikes ordinary oxygen molecules (O2) and disassociates them into single atoms of oxygen. The ultraviolet energy is absorbed by the oxygen and does not reach the ground. These single atoms are extremely reactive and combine with nearby oxygen molecules to form triatomic oxygen or ozone (O3)2.

The ultraviolet radiation that forms ozone also destroys it when the radiation strikes an ozone molecule and breaks it up into an ordinary oxygen molecule and a free atom of oxygen. Thus ozone in the stratosphere is constantly being created and destroyed so that it does not accumulate past a certain amount3. In my next post, I’ll discuss how ground-level ozone is formed.


  1. For example, see the Almanac of Policy Issues. To view, click here.
  2. NASA Earth Observatory, Chemistry in the Sunlight: The Chemistry of Ozone Formation To view, click here. You can see an excellent YouTube video of the process by clicking here.
  3. NASA Earth Observatory, Chemistry in the Sunlight: The Chemistry of Ozone Formation To view, click here.

EPA’s New Ozone Rule Part 2

Recently, the Obama administration withdrew a proposal to reduce the maximum allowable level of ground-level ozone concentration in the atmosphere1. The question that I wish to address is whether the benefits that might accrue to our nation from such a reduction are greater than the costs, particularly to industry. To analyze this problem, we need to understand what ground-level ozone is, how it is formed, what man-made processes promote ozone formation, and what industry must do to reduce the level of ozone.

Ozone is a form (called an allotrope2) of oxygen, the eighth element in the chemical periodic table3. Pure oxygen usually exists as molecules consisting of two oxygen atoms each, represented by the chemical formula O2. Ozone consists of molecules of three oxygen atoms each, represented by the chemical formula O3. Despite the fact that ozone consists of nothing but oxygen atoms, it is far more chemically reactive than ordinary oxygen4. For example, one cannot breathe pure ozone: breathing ozone in concentrations fifty parts per million or higher is probably fatal within 60 minutes5. Likewise, ozone can dissolve far more readily in water than ordinary oxygen6 and attacks substances (such as certain rubbers) that are not touched by ordinary oxygen7.

Breathing ozone is harmful to health even in low concentrations. Breathing air with 1.5 parts per million (ppm) of ozone for more than two hours can result in severe lung irritation with fluid-buildup, chest pain and cough, and extreme fatigue5. Ozone is known to attack and injure the tissues in the upper respiratory system, although the damage can be repaired by the body in a matter of weeks8.

It is important to distinguish between ozone in the troposphere (that part of the atmosphere that rests on the surface of the Earth) and the stratosphere (that layer of the atmosphere between about 6 and 31 miles above the surface at temperate latitudes). About 90% of all ozone in the atmosphere is in the stratosphere where it performs the very important function of absorbing high-energy ultraviolet radiation from the sun (all of the UV-c rays, most of the UV-b rays, and about half of the UV-a rays)9, preventing them from reaching the surface of the Earth where they would harm life. This ozone poses no dangers to humans; on the contrary, it helps make life possible. It is the 10% of the ozone in the atmosphere that exists in the troposphere (called tropospheric or ground-level ozone) that poses problems and is the subject of the proposed government regulation.

According to NASA, ground-level ozone levels without the presence of human activity should be about 10 to 15 parts per billion (ppb, one part per million equals 1000 ppb)10. Industrial activity has boosted those levels significantly such that the Environmental Protection Agency has established a limit of 80 ppb10. It appears to me that most people are able to breathe in that level of ozone without ill effects, or respiratory illnesses would be much more common than they are now. The question is whether people with respiratory problems, the very young, and the very old are adversely affected. If they are, is it cost effective to lower levels of ozone to improve their quality of life? Also, could long-term exposure to 80 ppb of ozone cause any significant health effects?

In my next post, I want to discuss how ozone is produced.


  1. Statement by the President on the Ozone National Ambient Air Qualities Standards. White House website. To view, click here.
  2. For a good explanation of allotropes, see the Diffen website, Oxygen vs Ozone.
  3. See the WebElements Periodic Table on oxygen.
  4. Rachel Cassiday and Regina Frey, Washington University.Chemical Properties of Ozone. To view, click here.
  5. Ozone Levels and Their Effects, edited by Den Rasplicka, OzoneLab Instruments website. To view, click here.
  6. Bruce Mattson, Janel Michels, Stephanie Gallagos, Creighton Univerisity.Microscale Gas Chemistry, Part 28 Mini-Ozone Generator: 800 nanomoles/minute p.7 paragraph “Office Paper.” To view, click here.
  7. Bassam Z. Shakhashiri, University of Wisconsin – Madison.Chemical of the Week: Ozone paragraph 7. To view, click here.
  8. U.S. Environmental Protection Agency website, Ground-Level Ozone: Health. To view, click here. For a more detailed treatment, see Health Effects of Ozone in the General Population, which you can view by clicking here.
  9. U.S. National Aeronautics and Space Administration Ozone Hole website, Ozone Facts tab, paragraph 2. To view, click here.
  10. Jeannie Allen, The Ozone We Breathe, NASA Earth Observatory website. To view, click here.