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.
Pingback: EPA’s New Ozone Rule: Part 24 | Michael Klein's Environmental Essays