1. Introduction.

Atmospheric chemistry is not a very old field, especially when compared to chemistry, physics, and meteorology. Even though carbon dioxide and oxygen were discovered over one hundred years ago, atmospheric chemistry as a field did not really get started until the 1950’s. Since then, progress in understanding atmospheric chemistry has been rapid, as researchers have applied the rapid advances in instrumentation, computers, and models and insights from other fields.

1.1 A timeline for atmospheric chemistry

A timeline for atmospheric chemistry shows this rapid progress (Table 1.1).

Table 1.1. A brief timeline of atmospheric chemistry

date	event or discovery	notes
1750	Joseph Black discovers carbon dioxide.
1774	Joseph Priestley discovers molecular oxygen.
1840	Christian Friedrich Shoenbein identifies ozone; J.L. Soret determines its OOO form 20 years later.
1858	Andre Houzeau makes first atmospheric ozone measurement.
1862	J.B. Bussingault shows that methane is in the atmosphere.
1881	W. N. Hartley shows that the atmosphere’s UV cut-off near 300 nm is due to stratospheric ozone	atm profile
1920s	Fabry, Buisson, and G.M.B. Dobson demonstrate that the ozone layer is ~3 mm thick and that it has latitudinal and seasonal variations	3 mm = 300 Dobson Units
1930	Sidney Chapman proposes mechanism to describe the abundance of stratospheric ozone. (The correct rate coefficients were not known for another 40 years.)
1940’s	Smog causes noticeable damage in Los Angeles
1952	London acid fog causes 4000 excess deaths; causing the UK to adopt pollution controls.
1950’s	Arie Haagen-Smith shows that sunlight+NOx+VOCs causes ozone, which is the most important component in LA smog
1950’s	F.E. Blacet discovers that photolysis of NO2 leads to O3
1963	Clean Air Act is first US environmental law; authorizes the EPA to create NAAQS; was amended in 1970, 1977, 1990.
1969-70	Julian Heicklen and others show that the hydroxyl radical, OH, drives daytime atmospheric chemistry
1971	Paul Crutzen proposes a stratospheric nitrogen mechanism that destroys ozone; Harold Johnston proposes that SSTs will destroy significant stratospheric ozone.	Nobel Prize, 1995
1974	Mario Molina and Sherwood Rowland propose that CFCs get into the stratosphere, are destroyed by UV sunlight, and the resulting chlorine destroys stratospheric ozone.	Nobel Prize, 1995
1985	Farman et al publish their measurements of the Antarctic Ozone Hole
1987	Anderson and coworkers demonstrate the link between Antarctic ozone loss and stratospheric chlorine; Montreal Protocol, the first international agreement to protect the atmosphere is signed; amended in 1990.

We can see from this timeline that progress was made simultaneously in tropospheric chemistry and stratospheric chemistry. In fact, advances in one part of the atmosphere have often led to advances in other parts of the atmosphere.

In some sense, atmospheric chemistry is a pure science, and in another sense, it is an applied science. One of the reasons for its rapid advance has been the pressing desire to develop a better understanding on atmospheric chemistry in order to mitigate or remove deleterious impacts that pollution has on humans, other animals, ecosystems, and structures. Studying the fundamentals of atmospheric chemistry has real relevance to us and our environment.

A prime example is the study of stratospheric ozone. From the mid 70’s, after Rowland and Molina had proposed their theory linking chlorofluorocarbons to stratospheric ozone destruction, a virtual army of researchers came from physics, chemistry, meteorology, the atmospheric sciences, engineering, biology, and computer sciences to study the problem. National and international programs were established, the issues were enumerated, satellites, airborne and ground-based instruments were built, models were developed and run, and meetings – many meetings - were held. By the mid 80’s, Rowland and Molina’s theory had been generally verified and many of the important issues had been resolved. This army of researchers began thinking that the stratosphere would be a solved scientific problem within a few years. Then the Antarctic ozone hole was discovered…

1.2 Impacts of atmospheric chemistry on us

Many of atmospheric chemistry’s issues – some very important to get right – are still unresolved (Table 1.2).

issue	cause	effects	uncertainty (my assessments)
Global warming	CO2, O3, aerosols	T, rain shifts, rising water levels	forcings – low to high; responses – low to high
Stratospheric ozone depletion	CFCs, HSCTs	Greater surface UV; cancer, plant damage	low to moderate
Urban ozone	VOC’s, NOx	Health, plant damage	moderate
Urban aerosols (PM2.5)	SO2, NOx, organics	Sudden heart attacks; asthma; low visibility	moderate
Regional air quality	same as urban	Same as urban; reduced air quality downwind	moderate
Global air quality	urban smog / biomass burning	Reduced air quality downwind; altered atmospheric oxidation potential; viruses	moderate to high
Acid rain	SO2, NOx	ecological damage; structural damage	low

Many of these issues are linked, so that while some aspects of each issue can be studied in isolation, others cannot. To make matters worse, strategies to mitigate environmental problems – even when based on sound scientific and technological principles – can easily fail if the human element has not been taken into account.

One example of this is Mexico City, a megalopolis of 24 million people. Air pollution is a serious health issue, and much of this pollution is related to the cars and trucks. The Mexico City Metropolitan Area government decided to restrict the use of every car to 5 days from Monday to Saturday by numbering the license plates and fining anyone who was driving on a day when the car was supposed to be idle. However, instead of increasing ridership on buses and the subway, it actually decreased it. Mexico City, which had been a net exporter of older automobiles to the rest of Mexico, became a net importer as residents bought a second car so that they could drive every day. These older cars polluted more. To make matters worse, the residents then had another car to drive on Sunday, when every car could be used. The flaws in this program have now been recognized and efforts are being made to modify it so that it will work.

1.3 Lifecycles of atmospheric constituents

Most atmospheric constituents either are emitted from Earth’s surface or a created by chemistry from other atmospheric constituents that are emitted from Earth’s surface. Some are quite stable in the atmosphere while others are very reactive. Yet ultimately, essentially all atmospheric constituents are returned, in one form or another, back to Earth’s surface.

Atmospheric constituents have a wide variety of sources, transformations, and sinks (Figure 1.1). Meteorological processes are as important in determining the evolution of atmospheric constituents as is the atmospheric chemistry. Emissions are natural (volcanoes, biogenic sources like trees, crops, and microbes, oceans, insects and wild animals) or anthropogenic (transportation, power generation, industry, agriculture, solvents, and other human products). These emissions are transported by turbulence, convection, clouds, weather systems and global wind patterns. Emissions can be transformed by atmospheric chemistry in the gas, liquid, or solid phases, often but not always in the presence of sunlight. Usually, chemicals evolve from forms that are more reduced to forms that are more oxidized, since the atmosphere is an oxidizing environment. More oxidized chemicals tend to be more water soluble and “stickier”. Sinks for atmospheric constituents includes uptake in biota, dry deposition on surfaces, uptake into cloud drops and subsequent rain out (wet deposition), and air-sea exchange.

The characteristic time that an atmospheric constituent remains in the atmosphere is often called its atmospheric lifetime, or residence time.

rate of accumulation of a constituent in a volume

=

rate of constituent coming in

-

rate of constituent going out

+

emission rate of constituent

-

removal rate of constituent

(1.1)

If Q is the total constituent mass (or number), F_in the mass flow rate in, F_out the mass flow rate out, P the emission (or production) rate, and R the removal (or loss) rate, then

(1.2)

In steady state, Q does not change with time and dQ/dt = 0. In this case F_in + P = F_out + R. The residence time is given by the expression:

(1.3)

In steady state, we can also write the residence time as the expression:

(1.4)

If we are interested in the entire atmosphere, then F_in and F_out both equal zero.

Very often, the removal rate of a constituent is proportional to its mass, R = k Q, then τ = 1/k. k is called the first-order removal rate and has units of s^-1. If two removal processes are both proportional to Q, then

(1.5)

If τ₁ >> τ₂, then the constituent residence time for the first process is much longer than that for the second process. The second process will be mainly responsible for the removal of the atmospheric constituent. We often use the concept of residence time, or lifetime, to tell us which processes are the most important for the atmospheric chemistry that we are studying.

Problem 1.1 The current mixing ratio of methane (CH₄) is ~1760 parts per billion by volume (ppbv) and is increasing at a rate of about 8 ppbv per year. Its main sinks are reaction with the hydroxyl radical (OH) in the atmosphere (k_OH = 3x10^-9 s^-1), and removal by soils (k_soils = 2x10^-10 s^-1).

a. What are the lifetimes (in years) of methane due to the two different processes?

b. What is the total methane lifetime (in years)?

c. Which process – reaction with OH or soils – is the most important in determining methane’s lifetime?

In reality, because methane partly controls the amount of atmospheric OH, the reduction of methane will cause shifts in atmospheric chemistry. Because of this feedback, the real amount of time required to remove all the methane, called the adjustment time, is 12 t o17 years for methane, different from the lifetime.

1.4 Units

Atmospheric chemistry and aerosol science have a large range of units. You will need to become adept at going among them.

Many of the units that are routinely used in atmospheric chemistry do not conform to MKS units. Some are cgs. Some are based on number of molecules. Some are based on mass.

Often, units originate from the original and most popular measurement techniques. (e.g, mass for aerosols).

Units fall into three classes:

1 mass or volume mixing ratio

a You must specify by mass (with an “m” at the end) or by volume (with a “v” at the end).

b c is often used as the mixing ratio, but not always.

c Common units are: ppm = 10^-6, ppb = 10^-9; ppt = 10^-12; ppq = 10^-15.

d The equivalent “correct” MKS units to volume mixing ratios is :

i) mmol mol^-1 = ppmv;

ii) nmol mol^-1 = ppbv;

iii) pmol mol^-1=pptv;

e Sometimes, the amount of a constituent will be designated as ppbC. Thus, for constituents that have more than one carbon, you would need to divide by the number of carbons in the constituent to get the actual mixing ratio in terms of molecules.

f Mixing ratios of gases and aerosols are transported.

2 mass, molecule number, or moles per unit volume (called concentration, or density)

a You must specify mass per unit volume or number of molecules per unit volume.

b n_A or [A] is often used to specify concentration.

c For gases, molecules cm^-3 is the most common unit.

d N_o/standard volume = 6.02x10²³/22,400 = 2.69x10¹⁹ molecules cm^-3 at STP.

e To go to other p and T, n_M = 2.69x10¹⁹ (p/1013) (273/T) = 7.25x10¹⁸ (p/T) molecules cm^-3.

f Concentrations of gases react.

3 mass or number of molecules per unit area integrated vertically between the observation point and space (called column density)

a important for remote sensing of atmospheric gases.

b to get this, must integrate the concentration of the gas from the surface to space: N = ò_z^¥ n_A dz

c usual units are molecules cm^-2 .

d column ozone is measured in Dobson Units (DU). 300 DU is equivalent to a 3 mm thick layer of pure ozone at STP (1013 hPa, 273 K).

4 conversions between mixing ratios and concentrations for constituent A

a For mixing ratio to number concentration, n = ppmv x 10^-6 n_M

b For mixing ratio to mass concentration, mg m^-3 = ppmv x (1000/22.4) (p/1013) (273/T) (molecular weight) = 12.03 (p (in hPa)/T (K)) M_A;

c For number concentration to mass concentration, mg m^-3 = molecule cm^-3 x 10¹² (M_A (in g))/N_o

You will need a periodic table to find atomic mass per mole. One on-line periodic table can be found at the website: http://pearl1.lanl.gov/periodic/default.htm .

Example. In a typical State College summer’s day, T ~ 80^o F, p = 960 hPa, and c_O3 ~ 80 ppbv. What is ozone’s number concentration and mass concentration? T = 26.7^o C = 300 K. n_M = 7.25x10¹⁸ (960/300) = 2.32x10¹⁹ cm^-3. Therefore for number concentration, n_A = 80 ppbv x 10^-9 x 2.32x10¹⁹ = 1.85x10¹² cm^-3. Ozone’s molecular mass is 3x16 = 48 g mol^-1. Therefore, the mass concentration is 0.080 x 12.03 x (960/300) x 48 = 148 mg m^-3.

Problem 1.2. On the same typical State College summer’s day, plumes from power plants in the Ohio River Valley can contain as much as 10 ppbv of sulfur dioxide (SO₂).

1. What is the number concentration of SO₂?
2. What is the mass concentration of SO₂?
3. If all of this SO₂ is converted over to sulfuric acid, H₂SO₄, what is the resulting mass per unit volume of the sulfuric acid?

Global numbers:

How many molecules are there in the atmosphere?

N_atm = òòò n_M dV ~ 4pR_E² ò_z n_M(z) dz

R_E = 6400 km = 6.4x10⁶ m = 6.4x10⁸ cm

What is the vertical distribution of n_M?

The first steps in this derivation are in Box 2.1 on page 27. Assuming a constant temperature, about 270 K, and doing the integration, we get

p = p_o exp(-z/(RT/Mg)) = p_o exp(-z/H), where H = RT/Mg = (8.31 270) / (0.029 9.8) ~ 7.5 to 8 km, is called the scale height because for z = H, p has dropped e(-1) from the surface value.

By the ideal gas law, p = n kT, where n is the number concentration, so n = n_o exp(-z/H), if we assume that T is constant.

Doing this integration, N_atm = 4pR_E² ò_z n_o exp(-z/H) dz = 4pR_E² n_o H = 4p (6.4x10⁸)² 2.5x10¹⁹ 7.8x10⁵ = 10⁴⁴ molecules

90% of the air is in the troposphere; 10% is in the stratosphere. So there are roughly 10⁴³ molecules in the troposphere.

What is the mass of the atmosphere? M_air = 0.029 kg mole^-1. So mass_atm = (10⁴⁴ / 6.02x10²³ ) 0.029 = 4.8x10¹⁸ kg.

Can relate this to pressure by dividing by Earth’s surface area, multiplying by g, and dividing by 100 to get hPa = mb.

Problem 1.3. The mean global stratospheric ozone column is about 300 DU.

1. What is the column density in molecules m^-2?

2. What is the total number of ozone molecules in Earth’s atmosphere?

3. What is the total mass of ozone in Earth’s atmosphere?

4. What percentage the ozone is lost each year in the Antarctic ozone hole?