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4.5 The Story of the Atmosphere's PAC-MAN

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    Hydroxyl (OH) is the PAC-MAN™ of the atmosphere. It reacts with thousands of different molecules, including carbon monoxide (CO), nitrogen dioxide (NO2), methane (CH4), sulfur dioxide (SO2), and ethane (C2H6).

    Credit: W. Brune

     

    The atmosphere's oxidation capacity is its ability to clean itself of all of the gases that are emitted into it. What does stratospheric ozone have to do with the atmosphere’s oxidation capacity, which mostly occurs in the troposphere and mostly by the atmosphere's PAC-MAN, hydroxyl (OH)? It turns out that natural dynamic processes actually pull air down from the stratosphere and mix it into the troposphere, eventually mixing some of this ozone to Earth’s surface. This naturally occurring surface ozone provides a baseline value for near-surface ozone, but ozone pollution is more than ten times greater than this baseline in cities. Ozone is both sticky on surfaces and fairly reactive in the atmosphere. It is lost both by depositing on surfaces and through being chemically destroyed by reactions in the atmosphere.

    The following chemical sequences are the humble beginnings of the atmosphere's PAC-MAN. OH is generated throughout the stratosphere and troposphere by a two-step reaction sequence. The first step is:

    \[O_{3}+U V \rightarrow O_{2}+O^{*}\]

    where O* is an excited-state oxygen atom that has extra chemical energy. O* can lose this extra energy by colliding with N2 and O2, but it can also collide with a water molecule to make two OH molecules:

    \[O^{*}+H_{2} O \rightarrow O H+O H\]

    OH is very reactive. You can think of OH as being water that has had a hydrogen taken away and wants it back. There are other sources for OH, but this one is the most important globally. OH reacts with many other atmospheric constituents. In fact, it is so reactive, that its lifetime in the atmosphere is less than a second.

    Another important oxidant is nitric oxide (NO). It comes from combustion (power plants, internal combustion engines, fires) or lightning. In cities, the NO mixing ratio is tens of ppb (10-9, by moles) during morning rush hour and a bit smaller during evening rush hour, but there is typically about a ppb around during the day. In very remote areas, the levels of NO are a hundred times less. NO can react with many other chemicals, but it reacts with O3:

    \[N O+O_{3} \rightarrow N O_{2}+O_{2}\]

    which forms nitrogen dioxide, NO2. NO2 is not very stable:

    \[N O_{2}+\operatorname{nearUV} \rightarrow N O+O\]

    ut the O reacts immediately with O2 to form ozone:

    \[O+O_{2}+N_{2} \rightarrow O_{3}+N_{2}\]

    If an NO2 molecule is produced, then an O3 molecule will be produced during the day when the sun is out. Note that if we think of these three reactions as a cycle, no ozone was either created or destroyed because it is destroyed in [4.5] and created in [4.2].

    What happens to all of the methane emitted into the atmosphere?

    Methane is a volatile organic compound (VOC). Methane oxidation is a good model for what happens to all of the volatile organic compounds that you smell every day and all the ones that you can’t smell. I am not going to show you the entire reaction sequence. Instead, here are just a few steps.

    The first step is the reaction between methane and hydroxyl:

    \[C H_{4}+O H \rightarrow C H_{3}+H_{2} O\]

    Note that water vapor is made and CH3 is a radical because it has 12+3 = 15 protons and, therefore, electrons. Just as for most other VOCs, and some other trace emissions, the reaction with OH is the main way methane is removed from the atmosphere. Otherwise, it would build up to high abundance.

    CH3 is very reactive. It adds an O2:

    \[\mathrm{CH}_{3}+\mathrm{O}_{2}+\mathrm{N}_{2} \rightarrow \mathrm{CH}_{3} \mathrm{O}_{2}+\mathrm{N}_{2}\]

    If there is any NO around, the following reaction happens:

    \[\mathrm{CH}_{3} \mathrm{O}_{2}+\mathrm{NO} \rightarrow \mathrm{CH}_{3} \mathrm{O}+\mathrm{NO}_{2}\]

    followed by:

    \[C H_{3} O+O_{2} \rightarrow C H_{2} O+H O_{2}\]

    and:

    \[H O_{2}+N O \rightarrow O H+N O_{2}\]

     

    The chemical CH2O is formaldehyde. Some of you may have encountered it in high school chemistry or biology and so may be familiar with the smell. You also see that we got the OH molecule back.

    Ultimately, formaldehyde gets broken down to CO and the net reaction of methane oxidation is:

    \[C H_{4}+2 N O+3 O_{2} \rightarrow \rightarrow \rightarrow \rightarrow C O_{2}+2 H_{2} O+2 N O_{2}\]

    Remember that NO2 is easily broken apart by the UV sunlight that reaches Earth’s surface, so we can take this reaction sequence a step further and show that in the presence of sunlight, reactions [4.6] and [4.2] give:

    \[N O_{2}+\operatorname{nearUV}+O_{2} \rightarrow N O+O_{3}\]

    or

    \[C H_{4}+5 O_{2} \rightarrow \rightarrow \rightarrow \rightarrow C O_{2}+2 H_{2} O+2 O_{3}\]

    In this final chemical equation, we do not see OH, HO2, NO, or NO2, yet they are essential to the formation of ozone. They are catalytic, which means that they are neither created nor destroyed in the reaction sequence, but instead are simply recycled between OH and HO2 and between NO and NO2. There are other reactions that destroy these reactive chemicals by producing other chemicals that are much less reactive and sticky, a main one being:

    \[O H+N O_{2}+N_{2} \rightarrow H N O_{3}+N_{2}\]

    where HNO3 is nitric acid, a very sticky and water soluble chemical. However, each OH that is produced can typically oxidize more than ten methane molecules before it reacts with NO2 to form nitric acid. And as reaction [4.13] shows, each time methane is completely oxidized, two O3 molecules are produced. That's a lot of chemical steps to remember, but I don't want you to necessarily remember them. I want you to see that the process started with a reaction of OH with a volatile organic compound (in this case methane) and that in the subsequent reactions, the product molecules had more and more oxygens attached to them. This process is why we say that the atmosphere is an oxidizing environment.

    Where does ozone pollution come from?

    Ozone is a different sort of pollutant from others because it is not directly emitted by a factory or power plant or vehicle but instead is produced by atmospheric chemistry.

    Three ingredients are needed to make ozone pollution: volatile organic compounds (VOCs) (like methane); nitric oxide (NO from combustion); and sunlight. When we say this, we assume that we already have some ozone and water to provide the OH to get the reactions started. Every VOC goes through an oxidation process that is similar to the methane oxidation reaction sequence. In the methane oxidation sequence, steps [4.9] and [4.11] make NO2, which in the presence of sunlight makes ozone through step [4.6] followed by step [4.2]. Voila! Ozone is formed from methane oxidation in the presence of nitrogen oxides and sunlight. Now imagine the thousands of volatile organic compounds in the atmosphere and realize that all of them - both anthropogenic and natural - can participate in the production of ozone pollution. Now you have seen the sequence of chemical reactions that produce tropospheric ozone.

    Let's look at a video (3:14) entitled "Ground Level Ozone: What Is It?" that explains ozone production without getting into the gory details of the chemistry.

    Ground Level Ozone: What is it?

    Click here for transcript of the Ground Level Ozone video.

    We're all pretty familiar with what O2 is. I hope so. You need to breathe it to live. Yes, O2 is oxygen, that life-giving gas, but what is O3? O3 is another gas essential to our survival but it's definitely not for breathing. O3 is ozone high up in the stratosphere. It's made naturally and absorbs harmful ultraviolet rays from the Sun. Without it life as we know it wouldn't, couldn't exist. We need the ozone layer in the stratosphere. We want it, we rely on it. But don't get too used to singing ozone's praises. High ozone levels at lower altitudes, what we call the troposphere, where we live and breathe or anything but natural and beneficial. In fact, down here it turns out to be a toxic atmospheric pollutants. Yep, you heard me right. ground-level ozone primarily exists due to human activities that burn fossil fuels. Transportation, our power and industrial plants, and other activities expel nitrogen oxides and hydrocarbons. When those compounds interact with sunlight, voila, ozone is created a contributor to smog. that's why I ozone levels increase during the summer months when sunlight is abundant. Yes, smog love summer just like many of us. We run, bike, hike, fish, play, stroll, oh yeah, and breathe. Yes, the fact that more people are outside when it's warmer makes us particularly vulnerable to Ozone's harmful impacts. Ozone is a harmful oxidant when we inhale it it's like getting a sunburn inside your lungs and it can be particularly serious for the young, old, active, and those with respiratory conditions at any age. And it's not just humans that are vulnerable ozone harms plants, crops, and agricultural yield interfering with pretty important processes like well, photosynthesis and even our economy. To make matters worse ozone production increases with higher temperatures which are occurring more frequently with climate change. The EPA sets national ambient air quality standards for several pollutants in the United States including ground level ozone. When a county is out of compliance they need to know what can be done to improve air quality. and let's not forget that air pollution is a global comments. air pollution is shared from surrounding cities states and also country's halfway around the world. What can we do, what are we willing to do to improve current levels? Drive less, carpool, avoid car idling, set your home's thermostat higher in the summer and lower in the winter, avoid gas powered lawn & garden tools on severe ozone days. There's a lot to do and lots to know about air quality knowing more about the sources and contributors to ozone and other atmospheric pollutants will help us chart our course.

    Credit: UCARConnect

     

    Ozone pollution is bad for the health of people, crops, and forests. Ozone can react with some types of VOCs, including types that make up our lungs, and breathing it can cause serious health problems and even death. Ozone reacts with the VOCs that make up plants and stunts their growth and damages their fruit. The Clean Air Act from the 1970s has dramatically decreased the levels of air pollution in the United States, including ozone. The EPA can take the credit for much of the progress against air pollution in the United States. But there is still a ways to go and the progress may be reversed due to effects of climate change. Since ozone pollution increases at higher temperatures, the increases in global temperatures could actually reverse the steady progress in ozone reduction and ozone pollution could once again increase, unless volatile organic compounds and nitrogen oxides are reduced even more.

    Now you can see why OH is called the PAC-MAN of the atmosphere. But how can we tell how long it will take for OH to remove from the atmosphere some trace gas like methane? Let’s look at an equation for the budget of methane. It is produced in the atmosphere by all the emissions from cows and wetlands. It is removed from the atmosphere by reactions with OH [4.7]. The rate of removal, that is the change in the methane concentration, is always proportional to the amount of the two reactants, in this case, CH4 and OH. So, the change in methane is given by the balance between the production and the loss by reaction with OH:

    \[\frac{d\left[C H_{4}\right]}{d t}=\) production \(-k_{O H+C H 4}[O H]\left[C H_{4}\right]\]

     

    where kOH+CH4 is the reaction rate coefficient (units: cm3 molecule-1 s-1) and [OH] and [CH4] are the concentrations of OH and CH4 (units: molecules cm-3).  Note that the production is positive and increases CH4 with time while the loss is negative and decreases CH4 with time.

    We use [OH] to indicate the concentration of OH (molecules cm-3), which is quite different from the OH mixing ratio (ppt = 10-12, or ppb = 10-9). 1 ppt ~ 2.4x107 molecules cm-3 and 1 ppb ~ 2.4x1010 molecules cm-3 for typical surface conditions. See the video below (1:47) entitled "Rate Equation" for further explanation:

    Rate Equation 

    Click here for transcript of the Rate Equation video.

    Let me explain equation 4.15, which is a rate equation for methane. A rate equation is just a differential equation. The change of something with respect to time equals the production rate of something, minus the fraction of something that is lost each unit of time, multiplied by the amount of something. Note that the loss rate of something is always proportional to something. That something can be anything. It does not have to be a chemical concentration. It could be the amount of milk in your refrigerator, or the number of socks in your drawer, both of which tend to disappear over time. And equation 4.15 is the methane concentration, which has units of molecules per centimeter cubed. The production rate is in units of molecules per centimeter cubed per second. Remember, each term of the equation must have the same units. The last term is the loss rate. The reaction rate coefficient has units of centimeter cubed per molecule per second, but when we multiply it by the OH concentration, we get a product that has units of per second, which is a frequency. Now, OH varies from almost 0 at night, to a peak value at midday. However, we can take an average OH to find the average loss rate of methane. Note that if we assume that the production rate suddenly goes to 0, then we find a very simple equation, which has an exponential solution. We designate the time that it takes the exponential factor to go to minus 1 as a lifetime, which is just the inverse of a loss frequency.

    How can we find out what the lifetime of methane is? We assume that the production suddenly stops and equals 0. Then [4.15] becomes:

    \[\frac{d\left[C H_{4}\right]}{d t}=-k_{O H+C H 4}[O H]\left[C H_{4}\right]\] 

    \[\frac{d\left[C H_{4}\right]}{\left[C H_{A}\right]}=-k_{O H+C H 4}[O H] d t\]

    kOH+CH4 is the reaction rate coefficient for this reaction. Assume that OH is constant. Because OH is generated mostly from sunlight, it follows the sunshine and is greatest near midday and is very small at night. However, we assume that the OH concentration is the average over the day and night in order to assign it a constant value. Now integrate both sides of the equation:

    \[\int_{\left[C H_{4}\right]_{0}}^{\left[C H_{4}\right]} \frac{d\left[C H_{4}\right]}{\left[C H_{4}\right]}=\int_{0}^{t}-k_{O H+C H_{4}}[O H] d t\]

    \[\ln \left(\left[C H_{4}\right]\right)-\ln \left(\left[C H_{4}\right]_{0}\right)=-k_{O H+C H_{4}} \overline{[O H]} t\]

    \[\ln \left(\frac{\left[C H_{4}\right]}{\left[C H_{4}\right]_{0}}\right)=-k_{O H+C H_{4}} \overline{[O H]} t\]

    take exponential of both sides

    \[\frac{\left[C H_{4}\right]}{\left[C H_{4}\right]_{0}}=e^{-k_{O H+C H_{4}}} \overline{[O H]} t\]

    \[\left[C H_{4}\right]=\left[C H_{4}\right]_{0} e^{-k_{O H+C H_{4}}} \overline{[O H]} t\]

     

    So we see that methane decreases exponentially with time.

    The atmospheric lifetime is defined as the time it takes something to decrease to e-1 = 0.37 of its initial value. So to find the lifetime of methane in the atmosphere, we see when kOH+CH4[OH] t = 1, or:

    \[\tau=\frac{1}{k_{O H+\mathrm{CH} 4}[O H]}\]

    where τ indicates the lifetime. kOH+CH4 =3x10-15 cm3 molecule-1 s-1 and [OH] ~ 106 molecules cm-3, so:

    \[\tau=\frac{1}{3 x 10^{-15} 10^{6}}=3 \times 10^{8} \mathrm{seconds} \sim 10\] years

    This reaction rate coefficient is fairly slow. Other VOCs have reaction rate coefficients that are typically hundreds to hundreds of thousands of times faster, so the lifetimes of most VOCs is hours to days.

    The atmospheric lifetime of a gas is very important for determining how far a gas can travel from its source. Some trace gases have lifetimes of hours, so unless they are made by atmospheric chemistry, they can't travel more than a few tens of kilometers from their sources. Other gases have much longer lifetimes - methane is a good example with its 10-year lifetime. In 10 years, it can travel from its sources to most anywhere around the globe, even to the stratosphere. NASA measures the amounts of several gases from space. An excellent NASA website for accessing these satellite data and having it plotted as global maps is the Center for Trace Gas Data & Information Website at the NASA Goddard Space Flight Center's Earth Sciences Distributed Active Archive Center (GES DISC).

    This concept of atmospheric lifetime is pretty important. For instance, what if an industry is spewing a chemical into the atmosphere that is toxic at a certain concentration in the atmosphere? Then it is important to know if that chemical is removed in less time than it takes to become toxic or if it is going to continue to build up at toxic levels and not leave the atmosphere for a long, long time. If its atmospheric lifetime is hundreds to thousands of years, then maybe we shouldn’t let that industry dump that chemical into the air.

    Now it is your turn to solve some problems concerning the atmospheric lifetimes and rates of change for atmospheric constituents in a quiz.

     

    Quiz 4-1: Atmospheric lifetimes calculations.

    1. Find Practice Quiz 4-1 in Canvas. You may complete this practice quiz as many times as you want. It is not graded, but it allows you to check your level of preparedness before taking the graded quiz.
    2. When you feel you are ready, take Quiz 4-1. You will be allowed to take this quiz only once. Good luck!