Atmospheric_chemistry by pengxuebo


									 Dry adiabatic lapse rate:
(Typically ~10°C/1000m).
Lapse rate = temperature change (cooling) that results from a given adiabatic (no heat
loss or gain) change in altitude.
Consider a dry parcel of air close to the ground….

The “dry adiabatic lapse rate” tell us how much cooling will occur.
The lapse rate is a hypothetical change in temperature with changing elevation, which can
be thought of as a line with slope ΔT/Δelev.
Comparing this to the measured temperature vs. altitude profiles allows one to assess the
vertical stability of the atmosphere:
  a) Stable Atmosphere: If T predicted from hypothetical “lapse rate”is < true
     temperature profile, then the air parcel will not rise.

 b) Unstable Atmosphere: If T predicted from hypothetical “lapse rate”is > true
    temperature profile, then the air parcel will continue to rise.
Moist adiabatic lapse rate As a parcel of air rises and cools, its water vapor will
condense out when it reaches conditions of water vapor saturation, releasing latent heat.
This energy must be considered in lapse rate estimates. The rate of temperature change
associated with a given change in elevation will decrease if a parcel of air becomes water
saturated, because the T drop resulting from the P drop is partially offset by the latent
heat released by water condensation.
So, it is intuitive that the “moist adiabatic lapse rate” (~6°C/1000m) is lower than the
dry adiabatic lapse rate (~10°C/1000m).

An example of how the vertical stability of the atmosphere influences the dispersal of
fumes from a smoke stack.

 Temperature inversions
What are atmospheric temperature inversions?
They represent conditions of strong atmosphere stratification that inhibits vertical mixing
of the atmosphere.
Why are inversions significant?
Increased residence time of air in the atmospheric boundary layer allows more
contaminants to build up and more time for secondary pollutants to form.
Inversions can occur in several ways. In a sense, the whole atmosphere is inverted by the warm
stratosphere, which floats atop the troposphere with relatively little mixing. An inversion can
form from the collision of a warm air mass (warm front) with a cold air mass (cold front). The
warm air mass overrides the cold air mass in the frontal area, producing the inversion.
 Radiation inversions: are likely to form in still air at night when the earth is no longer
receiving solar radiation. The air closest to the earth cools faster than the air higher in the
atmosphere, which remains warm, thus less dense. Furthermore, cooler surface air tends to flow
into valleys at night, where it is overlain by warmer, less dense air.
 Subsidence inversions, often accompanied by radiation inversions, can become very
widespread. These inversions can form in the vicinity of a surface high pressure area when high-
level air subsides to take the place of surface air blowing out of the high pressure zone. The
subsiding air is warmed as it compresses and can remain as a warm layer several hundred meters
above ground level.
 A marine inversion is produced during the summer months when cool air laden with
moisture from the ocean blows onshore and under warm, dry inland air.

The Tropopause and the Troposphere
Tropopause: the horizon above which water vapor from the surface Earth's hydrologic
cycle no longer travels freely due to a temperature inflection that causes significant H2O
crystallization to ice.
The atmosphere below this is called the troposphere. Most of what we think of as
weather is caused by air movements within the troposphere. This is also the part of the
atmosphere that is directly related to the hydrosphere. There is water above this horizon,
but most of this water is produced and destroyed along with other upper atmosphere
constituents through the action of photochemical reactions like the oxidation of methane
in the stratosphere.
                             CH4 + 2O2 →→ → CO2 + H2O

The composition of the global atmosphere
The lower atmosphere is composed of: constituents that are constant in composition
throughout and naturally variable components, such as
water vapor N and C species produced by biological and human activity on the earth's
surface, S species from volcanic and human sources O3 from photochemical reactions
and human activities, plus very trace organic molecules of both natural and human
origins and a number of inorganic intermediates of chemical reactions occurring in the

Notice the various sources and sinks of minor and trace gasses in the troposphere in the
table below:
Chemistry of the Atmosphere
Most compositional components of the atmosphere are in a constant state of flux (except
for inert gasses).
Many compounds are created and consumed in chemical reactions facilitated by the
presence of active particle surfaces and the presence of light from the sun.
Inputs of various gaseous forms of C, H, O, N and S at the base of the atmosphere vary
by geographic region and from processes with daily and seasonal time scales (both
physical and biological in origin). Also various pollutants are added to the atmosphere by
human activities on the ground and in the sky.
There are many classes of chemical reactions in the atmosphere, as summarized in the
table below.
Classification of the roles various chemical species play is difficult because they may do
different things at different times.

Approximate lifetimes of important atmospheric compounds are shown in table below:
There is a rich diversity of compounds that are formed and destroyed during
photochemical processes (reactions mediated by light) in the atmosphere. Some of the
chemical species involved in these reactions are very reactive and exist only at steady
state as intermediates of other reactants and products.
In the more rarefied upper atmosphere (where chemical entities have more space to move
around without bumping into their gaseous neighbors) these reactive species can exist for
longer periods of time and at higher concentrations.

Photochemistry – light and chemical bonds
Photochemical reactions involve the interaction of light with reactants to produce
product(s). To understand how and why this works, we need to think back to the nature
of electronic bonding in molecules.
Bonds are composed of electrons in molecular orbitals (simplistically thought of as the
geometric combination of atomic orbitals on the two atoms forming the bond). Two
atomic orbitals will combine to create two molecular orbitals; usually one is of lower
energy than the other in terms of electrons that fill them. Just like atomic orbitals, each
molecular orbital holds one or two electrons of opposite spin quantum number.
The electronic ground state has all bonding electrons in the lowest possible energy
configuration for that molecule.
The excited state: If two bonding orbitals exist with some energy (∆E) between them in a
certain molecule and one is filled with an electron pair, there is a certain wavelength (λ)
of light, , that can cause one of the electrons to "jump" to the higher energy orbital, such
that the ∆E = h/λ = hv, where is the frequency (1/λ).
Additional higher energy levels may exist also requiring some other value of
hv to cause the excitation.
Energetic light, particularly in the upper reaches of the atmosphere, causes molecular
The further up ones goes (in the atmosphere), the more exciting it gets.
Some excited state molecules simply decay back to the ground state by re-emitting that
absorbed light. Others break apart into atomic or multi-atomic fragments of the original
molecule. And others still want nothing more than to share their excitement with other
molecules or molecule fragments and enter into chemical reactions that would not be
possible from the ground state.
Unpaired electrons in atoms or molecules are called radicals. A molecule with two
unpaired electrons of the same spin in two different molecular orbitals is said to be in the
triplet state. A molecule with two unpaired electrons of the opposite spin in two different
molecular orbitalsis said to be in the singlet state.
Most ground state molecules have paired bonding electrons in the same orbital. Most
singlet and triplet state molecules are excited.
Molecular oxygen is an exception: It's ground state is a triplet state and its excited states
include the paired state and the singlet state.

Excited state molecules can undergo all sorts of reactions, many of which are categorized
Dissociation of molecules into fragments (atoms or multi-atomic) can create radicals,
many of which are very reactive. Radicals and excited state molecules are the back bone
of atmospheric photochemistry. The granddaddy radical of them all is the hydroxyl
radical, HO , It mediates a large number of reactions in the atmosphere. Other important
                              .                      .                             .
radicals are atomic oxygen O , hydroperoxyl HOO , and nitrogen monoxyl, NO











Surface reactions
Chemical reaction that take place on the surface of atmospheric particles (dust, ice
particles, sulfuric acid crystals, etc..) are also important.
Some of the reactions types involving particles are summarized here. Particles that
nucleate atmospheric water droplets and then dissolve contribute largely to the chemistry
of rain.
It is important to distinguish photochemical and surface reactions that occur above and
below the tropopause.
The reactions that occur above the tropopause involve more excited states and more
unusual and reactive radicals, due to the higher energy of solar radiation at that altitude.
By the time light filters into the troposphere, it has lost a great deal of energy causing
excitations in the upper atmosphere, and contributes much less to the general
"excitement" of the lower atmosphere.
Inorganic gaseous pollutants (CO, O3, S, N, Cl)

Virtually all metals are present in the atmosphere at low levels. Particulate matter
emitted from combustion of fossil fuels contains trace metals that were present in the
original fuel sample. The greatest health hazard is from aerosols that are smaller than 2.5
μm in diameter and contain lead, beryllium, mercury, cadmium and chromium. These
particles are removed from the atmosphere by settling out over time or by precipitation
events. Mercury is the only metal that exists as a gas and is therefore the only metal that
exists in a steady-state concentration in air. All other metals are emitted to the
atmosphere from natural or anthropogenic sources, and then removed by settling or
precipitation events.

Gaseous inorganic substances remain in the atmosphere for long residence times, and
would continue to build up much higher levels than observed today if it were not for the
gas phase reactions that convert these substances to water soluble species. Table 36.1 lists
the atmospheric forms for gaseous inorganic pollutants.

Table 36.1 Gaseous inorganic pollutants

Element              Atmospheric Forms

Oxygen               O3

Carbon               CO, CO2
Chlorine            Cl-, CFCs

Nitrogen            NH3,N2O,NO,NO2,N2O5

Sulfur              H2S,SO2,SO3


Although ozone (O3) is a desirable substance in the stratosphere, it is a major
environmental hazard at ground level. Ozone is a by-product of photochemical smog and
reacts with hydrocarbons to form peroxynitrates that damage sensitive biological tissues
in the eyes, nasal passages, throat and lungs. Excessive ozone levels in the troposphere
have been blamed for killing plants through reactions with chlorophyll. Ozone is formed
naturally when oxygen molecules are photochemically dissociated into oxygen atoms that
can then react with a second oxygen molecule to make ozone. The presence of nitrogen
oxides (NO, NO2) leads to higher than normal background levels of ozone through
several well understood photochemical reactions.

Carbon Monoxide and Carbon Dioxide

Carbon monoxide is a product of the incomplete combustion of fossil fuels, and is often
listed in urban air quality measures. As much as 20% of the CO released to the
atmosphere each year comes from natural sources, but the greatest health problem is in
metropolitan areas near high densities of vehicular traffic. Urban air may contain carbon
monoxide at concentrations in excess of 100 ppm during rush-hour traffic. Catalytic
converters and air pumps have been installed on all motor vehicles built since the early
1970s as a way to reduce carbon monoxide emissions. Carbon monoxide is transformed
to carbon dioxide when exhaust gases are combined with air and passed over the surface
of a catalyst. Carbon monoxide has a 4-month lifetime in the atmosphere, where it reacts
with the hydroxyl radical to form carbon dioxide. The sequence of reactions that
transform carbon monoxide to carbon dioxide and regenerate the hydroxyl radical are
shown below:
Atmospheric carbon dioxide levels are determined by a long-term equilibrium between
CO2 in the air and CO2 dissolved in the oceans and surface water, releases of CO2 from
natural and anthropogenic sources, and losses by plant growth. The pre-industrial level of
carbon dioxide remained constant over the past 10,000 years, but has increased from 250
ppm in 1900 to about 350 ppm today. Since 1970, atmospheric CO2 levels have been
increasing at an exponential rate. Elevated levels of atmospheric carbon dioxide may
have a profound impact on the earth’s climate, and an international agreement that was
formulated in December 1997 in Kyoto Japan, attempts to reduce carbon emissions to
pre-1990 levels.


Chloride is water-soluble and washes out of the atmosphere quickly. Inorganic chlorides,
therefore, have very short lifetimes and do not constitute an environmental threat.

Chlorofluorocarbons (CFCs) and some volatile organic carbon substances (VOCs)
contain covalently bonded chlorine atoms, and being relatively inert substances, will to
remain in the atmosphere for several dozen years. By remaining in the atmosphere for
such long times, these molecules migrate across the troposphere/stratosphere boundary
where they become exposed to much higher levels of ultraviolet radiation than
experienced at ground level. Ultraviolet radiation can cause covalent bonds to separate
into individual atoms containing unpaired electrons. Any covalent bond can be broken if
sufficient energy is applied, and the ambient energy available in the stratosphere breaks
up larger molecules into individual atoms as if they had been placed on an anvil and
struck with a hammer. Molecules that contain chlorine produce chlorine atoms in the
stratosphere, and having an odd number of electrons, the chlorine atom results in a stable
free radical.
Sulfur Dioxide
Coal-fired electric power generating plants account for the majority of sulfur dioxide
emissions. Scrubbing technologies have been used since the early 1970s to remove SO2
from power plant emissions. This process is formally called "Flue Gas Desulfurization"
or FGD. All current technologies involve exposing the combustion gases to a substance
that will absorb most of the SO2. FGDs are catagorized into "wet" technologies that
expose the flue gases to an aqueous solution, and "dry" technologies that expose the flue
gas to solid absorbents. Wet technologies are more efficient at removing SO2 (~95 to
98% efficient), while dry technologies have slightly less efficiency (by a few percent) but
are considerably easier to handle. The chemical reactions for the most common FGDs are
given in the table below, along with comments about the advantages and disadvantages of
Table: Chemical reactions for the most common flue gas desulfurization technologies.

Process                       Reaction                                                 Notes

Lime slurry scrubbing         Ca(OH)2 + SO2 → CaSO3 + H2O                               1

Limestone slurry scrubbing CaCO3 + SO2 → CaSO3 + CO2(g)                                 2

Magnesium oxide
                              Mg(OH)2 + SO2 → MgSO3 + H2O                               3

                              Na2SO3 + H2O + SO2 → 2 NaHSO3
Sodium-base scrubbing                                                                   4
                              2 NaHSO3 + Heat → Na2SO3 + H2O + SO2

                              2 NaOH + SO2 → Na2SO3 + H2O
Double alkali                                                                           5
                              Ca(OH)2 + Na2SO3 → 2 NaOH + CaSO3(s)

    Up to 200 Kg of lime are needed per metric ton of coal, producing huge quantities of

2   Lower pH than lime slurry, not as efficient.

3   The sorbent can be regenerated, which can be done off site if desired.

4   No major technical limitations. Relatively high annual costs.

5   Allows for regeneration of expensive sodium alkali solution with inexpensive lime.

The lime slurry, limestone slurry and double alkali FGDs all have the advantage of
producing gypsum by oxidizing the calcium sulfite formed:

CaSO3 + ½ O2 2 H2O → CaSO4·2 H2O(s)
Gypsum has a small commercial value. The sodium base scrubbing process is used when
SO2 is recovered to manufacture sulfuric acid, another product with commercial value.

Acid Rain
What causes acid rain?
Acid rain is caused by the burning of fossil fuels. Burning oil, gas and coal in power
stations releases sulphur dioxide (SO2) into the atmosphere and burning oil and gas in
motor vehicles puts nitrogen oxides (NOX) into the atmosphere. These gases mix with
water droplets in the atmosphere creating weak solutions of nitric and sulphuric acids.
When precipitation occurs these solutions fall as acid rain.

Sulphur dioxide (SO2): is generally a byproduct of industrial processes and burning of
fossil fuels. Ore smelting, coal-fired power generators and natural gas processing are the
main contributors.
Nitrogen oxides NOx: the main source of NOx emissions is the combustion of fuels in
motor vehicles, residential and commercial furnaces, industrial and electrical-utility
boilers and engines, and other equipments.
How Do We Measure Acid Rain?
Acid rain is measured using a scale called "pH". The lower a substance's pH, the more
acidic it is. Pure water has a pH of 7.0. Normal rain is slightly acidic because carbon
dioxide dissolves into it, so it has a pH of about 5.5
What problems are caused by acid rain?
Acid rain increases the acidity levels of rivers, lakes and seas. This can kill aquatic life.
Acid rain increases the acidity levels of soils and this can kill vegetation. Acid rain has
been found to destroy the roots and leaves of forests in Germany and Scandinavia have
been destroyed as the result of acid rain emissions from the UK. Acid rain can erode
buildings and monuments (particularly if they are made from limestone).

Power Plants: Most sulfur dioxide comes from power plants that use coal as their fuel.
These plants emit about 100 million tons of sulfur dioxide, 70% of that in the world.

Automobiles: produce about half of the world's nitrogen oxide. As the number of
automobiles in use increases, so does the amount of acid rain. Power plants that burn
fossil fuels also contribute significantly to nitrogen oxide emission.
Natural Causes: fires, volcanic eruptions, bacterial decomposition, and lightening also
greatly increase the amount of nitrogen oxide on the planet. However, even the gigantic
explosion of Mt. St. Helens released onlyabout what one coal power plant emits in a year.

Industry can reduce acid rain creation by using coal with low sulfur content, they can
remove the sulfur from smoke their plants release, and they can limit processes known to
generate high levels of acid rain.
Environmentalists advocate the installation of sulfur cleaning scrubbers in factories,
washing sulfur out of coal, and finding new methods of burning coal. Power plant
operators are looking for less expensive solutions to the problem.
Photochemical Smog
What is photochemical smog?
Photochemical smog is a mixture of pollutants that are formed when nitrogen oxides and
volatile organic compounds (VOCs) react to sunlight, creating a brown haze above cities.
It tends to occur more often in summer, because that is when we have the most sunlight.
Primary pollutants
The two major primary pollutants, NOx and VOCs, combine to change in sunlight in a
series of chemical reactions, outlined below, to create what are known as secondary
Secondary pollutants
The secondary pollutant that causes the most concern is the ozone that forms at ground
level. While ozone is produced naturally in the upper atmosphere, it is a dangerous
substance when found at ground level. Many other hazardous substances are also formed,
such as peroxyacetyl nitrate (PAN).
What are the major sources of photochemical smog?
While nitrogen oxides and VOCs are produced biogenically (in nature), there are also
major anthropogenic (man-made) emissions of both. Natural emissions tend to be spread
over large areas, reducing their effects, but man-made emissions tend to be concentrated
close to their source, such as a city.
Biogenic sources
In nature, bushfires, lightning and the microbial processes that occur in soil generate
nitrogen oxides. VOCs are produced from the evaporation of naturally-occurring
compounds, such as terpenes, which are the hydrocarbons in oils that make them burn.
Eucalypts have also been found to release significant amounts of these compounds.
Anthropogenic sources
Nitrogen oxides are produced mainly from the combustion of fossil fuels, particularly in
power stations and motor vehicles. VOCs are formed from the incomplete combustion of
fossil fuels, from the evaporation of solvents and fuels, and from burning plant matter
(such as backyard burning) and wood-burning stoves.
How is smog formed?
Below is a simplified explanation of the chemistry of smog formation.
Nitrogen dioxide (NO2) can be broken down by sunlight to form nitric oxide (NO) and an
oxygen radical (O.):

Oxygen radicals can then react with atmospheric oxygen (O2) to form ozone (O3):

Ozone is consumed by nitric oxide to produce nitrogen dioxide and oxygen:
Harmful products, such as PAN, are produced by reactions of nitrogen dioxide with
various hydrocarbons (R), which are compounds made from carbon, hydrogen and other

The main source of these hydrocarbons is the VOCs. Similarly, oxygenated organic and
inorganic compounds (ROx) react with nitric oxide to produce more nitrogen oxides:

The significance of the presence of the VOCs in these last two reactions is paramount.
Ozone is normally consumed by nitric oxide, as in reaction 3. However, when VOCs are
present, nitric oxide and nitrogen dioxide are consumed as in reactions 4 and 5, allowing
the build up of ground level ozone.

How location and weather can have an effect
Topography: the topography of the area surrounding a city can vastly influence the
formation of photochemical smog. Because of the restriction of air movement, a city in a
valley can experience problems that a city on an open plain may not.
Meteorology: Normally the layer of air closest to the earth’s surface is warmer than the
air higher in the atmosphere because the heat of the sun is re-radiated (warmed by the
earth’s surface). The higher level cool air sinks and is then warmed and displaced
upwards in a convection cycle (Figure 1). This condition is called ‘unstable’ and helps to
carry pollutants upwards, where they are dispersed and diluted. This cycle is usually
assisted by higher wind speeds. However, when the opposite occurs - a temperature
inversion (cities can experience prolonged periods of photochemical smog).
An inversion is formed when a ceiling of warmer air traps the cooler layer of air, which
contains the pollutants, near the ground’s surface. This hinders the ability of the
pollutants to rise to the atmosphere and be dispersed. After an inversion has formed, it
keeps any smog that is present close to the ground, maximising its detrimental effect.
There are two major processes that enable an inversion to happen and both are usually
accompanied by low wind speeds.
The first, advection, is when an upper layer of warmer air is blown in, trapping the layer
of cool air below it. This ‘stable’ condition may last for several days. A variation of this
is when a cooler layer of air, such as a sea breeze, is blown in underneath a warmer layer,
creating the same effect.
The second process, radiation inversion, usually occurs overnight. The ground cools and
in turn cools the air layer closest to it, resulting in the lower air layer being cooler than air
above it, forming an inversion. (See figure below)
What are the dangers of photochemical smog?
Photochemical smog can have an effect on the environment, on people’s health and even
on various materials. The main visible effect is the brown haze that can be seen above
many cities. The brown tinge is caused by very small liquid and solid particles scattering
the light.
Chemicals such as nitrogen oxides, ozone and peroxyacetyl nitrate (PAN) can have
harmful effects on plants. These substances can reduce or even stop growth in plants by
reducing photosynthesis. Ozone, even in small quantities, can achieve this, but PAN is
even more toxic to plants than ozone.
The biggest concern about photochemical smog is the effect it has on people’s health.
The effects of the major primary and secondary pollutants in smog are given in table:
Table: Health effects of pollutants involved in photochemical smog

Ozone can damage various compounds. It can cause the cracking of rubber, the reduction
in tensile strength of textiles, fading of dyed fibers and cracking of paint.
How can we reduce the occurrence of photochemical smog?
The most effective way of reducing the amount of secondary pollutants created in the air
is to reduce emissions of both primary pollutants.
 Reduction of nitrogen oxide
The main method of lowering the levels of nitrogen oxides is by a process called
‘catalytic reduction’, which is used in industry and in motor vehicles. For example, a
catalytic converter fitted to a car’s exhaust system will convert much of the nitric oxide
from the engine exhaust gases to nitrogen and oxygen. In Australia, all motor vehicles
built after 1985 must be fitted with catalytic converters.
Nitrogen is not in the actual fuels used in motor vehicles or power stations; it is
introduced from the air when combustion occurs. Using less air in combustion can reduce
emissions of nitrogen oxides.
Temperature also has an effect on emissions—the lower the temperature of combustion,
the lower the production of nitrogen oxides. Temperatures can be lowered by using
processes such as two-stage combustion and flue gas recirculation, water injection, or by
modifying the design of the burner.
- Catalytic Converters
Catalytic converters on motor vehicle exhausts are a way of trying to reduce the carbon
monoxide and nitrogen oxide emissions. The catalyst used is either platinum or a
combination of platinum and rhodium.

The platinum catalyses the reaction of unburnt hydrocarbon (such as pentane) and oxygen
(O2) to produce carbon dioxide (CO2) and water vapour (H2O):

              C5H12                platinum         catalyst
                           + 8O2                             5CO2 + 6H2O
              (pentane)            ------------------->

The rhodium catalyses the reaction of carbon monoxide (CO) and nitric oxide (NO) to
form carbon dioxide (CO2) and nitrogen gas (N2):

                                    rhodium catalyst
                      2CO + 2NO                          2CO2 + N2

The reduction of nitric oxide (NO) to nitrogen gas (N2) must proceed more quickly than
the oxidation of carbon monoxide (CO) to carbon dioxide (CO2) or else all the carbon
monoxide will be oxidised to carbon dioxide before it can be used to reduce the nitric
oxide. Motor vehicles can only use catalytic converters if they use unleaded petrol since
the lead in petrol renders the catalyst inactive.
 Reduction of VOCs
There are various ways to reduce VOC emissions from motor vehicles. These include the
use of liquefied petroleum gas (LPG) or compressed natural gas (CNG) rather than petrol,
decreasing distances vehicles travel by using other modes of transport, such as buses and
bikes, and implementing various engine and emission controls now being developed by

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