Environment of Earth

September 16, 2009


Filed under: Atmospheric chemistry — gargpk @ 3:19 pm
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Most of the particulate material suspended in the  atmosphere has very small size and so has a very large surface area per unit mass (around 1 million square meter  per  gram). Such large surface area offers considerable opportunity for the absorption of molecules from the gas phase. This is particu­larly true if these molecules have a low volatility. A sub­stance having vapor pressure less than 10-6 Pa at ambient temperature will largely be adsorbed on the aerosol  particles. Therefore, metals volatilized through volcanic or biological processes will probably end up at­tached  to aerosols. The likelihood of surface reactions  also increased by the large surface to volume ratio of aero­sols. Generally,  two types of reactions occur  on aerosol: thermal reactions and photochemical reactions.

Thermal reactions: For describing thermal reactions on  aerosol surfaces, following two surfaces have been common  models of atmospheric aerosols:

(i) Sulfuric acid surface: Sulfuric acid is a liquid surface but acid covers the surface of many atmospheric aerosol particles so this is a good model. The effectiveness of sulfuric acid surfaces as sink has been investigated for a number of atmospheric trace gases. The effectiveness of surface may be measured in terms of the probability of reactions occurring on collision of the molecules of the gas with the surface. Such probabilities for some major atmospheric trace gases are given in the Table.

Table: Probabilities of reactions on collision of gas molecules with surface.

Molecule Probability
Water vapor 2 x 10-3
Ammonia >1 x 10-3
Hydrogen peroxide 7.8 x 10-4
Nitric acid 2.4 x 10-4

For  species like nitric acid or hydrogen peroxide, the absorption of the gas by sulfuric acid surfaces could be a sink of atmospheric gases as much important as the photolysis.

(ii) Graphite carbon surface: Absorption of gases by graphite carbon  is well known. A gas like sulfur dioxide is readily absorbed  and  presumably oxidized on the surface. However, aerosol  surface soon becomes saturated or poisoned.  Absorption  of gas molecules can not occur further unless there is some mechanism for ‘cleaning’ the surface. Thus it is  diffi­cult  to  visualize  the mechanism of the  removal  of  large amounts of a gas like sulfur dioxide from atmosphere by  such a heterogeneous solid phase process.

Photochemical reactions: In addition to possibility of  ther­mal  reactions on particle surface subsequent to the  absorp­tion  of the gas molecules, photochemical reactions are also possible. For example,


2CO + O2 —————-> 2CO2

TiO2, ZnO


2N2 + 6H2O ————-> 4NH3 + 3O2


The  importance of these reactions in the atmosphere  is  not known. However, it is known that photo-assisted reactions on titanium oxide or zinc oxide desert sands lead to  production of  ammonia. It has been postulated that such reactions were the source of ammonia in the early atmosphere of Earth.


Filed under: Atmospheric chemistry — gargpk @ 2:40 pm
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Most of the ozone in the atmosphere forms the Ozone layer in the stratosphere at altitudes between 10 and 40 km (100 to 0.1 mb pressure altitude) depending on latitude, just above the tropopause. This layer is crucial for life because only ozone absorbs UV-B radiation between 280-320 nm. UV-A rays between 320 and 400 nm are not affected by ozone while UV-C rays between 200 and 280 nm are absorbed by other atmospheric constituents also beside ozone.

Stratospheric ozone distributions are strongly dependent on stratospheric circulation patterns, varying according to latitude, seasons, short-term meteorological changes and the photochemical processes of formation and destruction. Major driving forces are availability of sunlight and thus of UV radiation and in upper stratosphere (above pressure altitude of 5 mb) the latitudinal temperature gradient which assists ozone transport. The ozone content of stratosphere is highly dynamic and variable. Its concentrations peak around the altitude of 30 km in tropics and around 15 to 20 km in polar regions.

Though hundreds of reactions are known to be involved in the ozone chemistry of stratosphere, only a few can be described properly. The ozone chemistry basically involves two types of reactions: those involved with ozone formation and those involved with ozone destruction. These two types of reactions are important because relationship between stratospheric ozone and climate has been studied particularly in association with ozone depletion and ultra-violet radiation. Another important feature is that above tropopause, liquid water does not play significant role and stratospheric ozone chemistry here is dominated by photochemical reactions.

1. Ozone formation: This itself is a photochemical process involving UV radiation of wavelength less than 242 nm. Though photodissociation of oxygen by UV radiation at less than 175 nm may yield an oxygen atom in excited state i.e. O(1D), such photodissociation is important only in the upper stratosphere because such short wavelength can not penetrate lower into stratosphere.

Thus in upper stratosphere reaction may be:

O2 + hv (<175 nm) ——> O(3P) + O(1D)

Oxygen atom in excited state on collision with some diatomic molecule (M2) yields oxygen atom in ground state i.e. O(3P):


O(1D) + M2 —————–> O(3P) + M2


while in lower stratosphere reaction is:

O2 + hv (175-242 nm) ———> O(3P) + O(3P)

The oxygen atoms in ground state react with diatomic oxygen molecules to form ozone:

O(3P) + O2 ———> O3

2. Ozone destruction: This involves those reactions which balance the photochemical formation of ozone in stratosphere:

O3 + hv ——> O2 + O(1D)

O3 + O ——-> 2O2

Another additional reaction for removal for oxygen atoms is:

O + O + M ——–> O2 + M

Many analogous reactions involving H, N and Cl radicals also occur in stratosphere:

OH + O3 ———> O2 + HO2

HO2 + O ———> OH + O2

NO + O3 ——–> O2 + NO2

NO2 + O ——–> NO + O2

O3 + Cl ——> O2 + ClO

ClO + O ——> O2 + Cl

All the above pairs of reactions are summed as:

O3 + O —–> 2O2

i.e. each pair of reactions involves destruction of ozone and atomic oxygen while restoring the OH, NO or Cl radical.


Filed under: Atmospheric chemistry — gargpk @ 2:35 pm
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N2O which is relatively stable in the troposphere, usually moves into stratosphere and undergoes following photochemical reactions:

N2O + hv (<260 nm) ————–> N2 + O(1D)

N2O + O(1D) ———-> N2 + O2

N2O + O(1D) ———-> 2NO

NO or NO2 which may also move from troposhpere into stratosphere or are produced in the stratosphere, undergo following reactions there:

NO + O + M ——> NO2 + M

NO2 + OH + M —–> HNO3 + M

HNO3 + hv (<330 nm) ——> OH + NO2

 Ozone layer in stratosphere absorbs sufficient amount of UV radiation so that at tropopause HNO3 has photochemical lifetime of about 10 days. This time is long enough for much of it to cross the tropopause and come down with rainfall thus being removed from stratosphere.


Filed under: Atmospheric chemistry — gargpk @ 2:30 pm
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Chief natural source of chlorine in atmosphere is probably methyl chloride from marine algae but it accounts for only 25% of the chlorine currently being transported across tropopause into the stratosphere. Other natural sources adding minor amounts are HCl acid from volcanoes and chlorine from sea sprays. In the past few decades, chlorofluorocarbons (mostly CFCl3 i.e. Feron-11) and CF2Cl2 i.e. Feron-12) added to the atmosphere by human activities have become chief source of stratospheric chlorine. Ammonium perchlorate-aluminium solid rocket propellents are another anthropogenic source of atmospheric chlorine. These compounds absorb UV radiation in the range of 190 to 220 nm resulting in their photodissociation:

CH3Cl + hv —–> CH3 + Cl

CFCl3 + hv ——> CFCl2 + Cl

CF2Cl2 + hv —–> CF2Cl + Cl

Free Cl atoms in stratosphere may undergo various reaction cycles:

1. Reaction with ozone: Free chlorine atoms in stratosphere react with ozone in catalytic manner and cause depletion of ozone:

Cl + O3 —–> ClO + O2

ClO produced may react with nitrogen compounds:

ClO + NO —–> Cl + NO2

ClO + NO2 + M ——–> ClNO3 + M

ClNO3 may be decomposed by UV radiation or by reaction with atomic oxygen:

ClNO3 + hv —–> ClO + NO2

ClNO3 + O —-> O2 + ClO + NO

Reactions of ClO with NO or NO2 are important because they effectively remove N- and Cl- species involved in ozone destroying cycles.

2. Reaction with methane and hydrogen: Free Cl may also react with CH4 or H2:

Cl + CH4 —–> HCl + CH3

Cl + H2 ——> HCl + H

Some of the HCl may react with OH radical in stratosphere:

OH + HCl —–> H2O + Cl

However, most of the HCl moves down to tropopause and is removed with rainfall as HCl acid.

September 13, 2009


Filed under: Atmospheric chemistry — gargpk @ 2:55 pm
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The ozone present in troposphere and stratosphere together constitute the total atmospheric ozone. The atmospheric ozone has important impact on the global climate system. The production and loss of ozone in both troposphere and stratosphere are strongly linked to atmospheric chemistry at both levels. Both areas of ozone are also influenced by four major processes that basically dominate the biogeochemical cycles in atmosphere:

1. Emissions from natural and anthropogenic sources

2. Chemical transformations and reactions

3. Atmospheric transport through circulation

4. Removal mechanisms

Ozone chemistry of troposphere

The troposheric ozone concentrations make up only 13% of total ozone in atmosphere yet ozone of this zone has major impact on climatic change through its effect on global warming. Natural background ozone concentrations can only be found in atmospheres of rural and remote areas while over urban centres, unnatural ozone concentrations are created by anthropogenic emissions of various substances that have profound effect on ozone chemistry. In unperturbed troposphere, the formation and destruction of ozone are part of a dynamic balance controlled mainly through ozone sources from marine and terrestrial biospheres and sinks atmospheric photochemistry and surface depositions. Anthropogenic emissions entering this system change the balance both spatially and temporally and such changes can be transferred globally by atmospheric transport mechanisms.

Naturally the tropospheric ozone is a secondary constituent originating from two main sources:

1. In upper troposphere, major source is transport of ozone from stratosphere

2. In middle and lower troposphere, photochemical mechanisms of ozone production

The concentration of ozone at any level in troposphere is determined mainly by photochemical mechanisms of its formation and destruction. Photochemistry dominates the ozone cycle particularly in middle and lower troposphere atleast for three reasons:

(a) Presently calculated rate of loss of ozone are about four times higher than the rate which would have occurred if tropospheric ozone originated completely in stratosphere

(b) Measured increase in ozone over urban areas can only be photochemical in origin

(c) Larger concentration of ozone in Northern Hemisphere than in Southern Hemisphere despite larger land surface sink can only be attributed to atmospheric photochemical reactions.

Recent estimates show that maximum ozone produced per year in troposphere is about 6.5 x 1011 molecules cm-1 s-1. Higher concentrations of ozone occur in mid-latitudes of Northern Hemisphere because of higher number of precurssor sources there. Minimum ozone concentrations occur in equatorial regions around 100 S caused partially by stronger photochemical destruction in the tropics and partially by background ocean conditions in Southern Hemisphere. Average ozone concentrations in free troposphere are 39 ppbv in Northern Hemisphere and 24 ppbv in Southern Hemisphere. Representative latitudinal ozone concentrations in free troposphere are 30-40 ppbv in 30-600 S, 20 ppbv in 0-300 S, 20-30 ppbv in 0-200 N and 30-50 ppbv in 20-600 N.

Vertical distribution of ozone differs between hemispheres and with distribution of important chemical precursors, particularly CO. In Northern Hemisphere, on average the ozone concentration increases slightly with altitude and boundary layer ozone concentrations are about 1.1 to 1.5 lower than the free troposphere. In Southern Hemisphere, there is little variation in ozone concentration with altitude and ozone in boundary layer does not decrease significantly compared to free troposphere.

The formation and destruction of ozone in troposphere depends heavily on the OH radical concentration and associated reaction efficiency. The process is initiated by photodissociation of ozone by sunlight and the formation of OH from water and oxygen:

O3 + hv (300-330 nm) ——-> O2 + O(1D)

O(1D) + H2O ——> OH + OH ( R = 2.3 x 10-10)

OH formation depends on water in troposphere. As a rough estimate, H2O0.5-1.0 approximates OH concentrations. Since Oh is a highly reactive radical, it is very short-lived in troposphere. Its concentrations sow diurnal variations, particularly in higher latitudes linked to solar-energy variations. At night, OH concentrations are supposed to fall by two orders of magnitude as compared to daytime with minimum concentrations about 105 molecules cm-3 and maximum concentrations near mid-noon about 107 molecules cm-3.

In the troposphere, apart from OH radical other critical species for basic gas-phase reactions are nitrogen oxides (NO, NO2, Nox), free hydrogen/oxygen radicals (OH, HO2), methane and non-methane hydrocarbons (designated by general term RO2 and carbon monoxide. These processes are strongly linked to one another and depend heavily in the concentrations of the relevant molecules in the atmosphere. These reactions in troposhpere have been described in detail in the discussion of photochemical smog problem. However, important features of main molecules affecting tropospheric ozone may be summarized as following:

1. Nitrogen oxides: Nitrogen gases help control OH concentrations in troposphere and concentrations of NO and NO2 are needed to form ozone. Since both molecules are active in ozone process, they are described by their conserved quantity, NOx. The rate of ozone production in troposphere seems to be controlled by NOx concentration. NOx acts as catalyst to photochemical reaction processes and provides the environment which allows further ozone formation or loss reactions in various chains. For example, in NOx-poor environment, oxidation of one methane molecule to carbon dioxide via CO results in net loss of about 3.5 H atoms and 1.7 ozone molecules. In NOx-rich environments, the same process will create about 0.5 H atoms and 3.7 ozone molecules. The transfer point between ozone loss and ozone production seems to an NO concentration of about 30.0 pptv. The efficiency of NOx in ozone-formation processes decreases with increasing NOx concentration. However, in terms of total production of ozone, this inefficiency is overcome in the atmosphere with higher NOx.

Main sink of NOx in atmosphere is conversion to nitric acid by OH. This sink acts within a time frame of 1 to 2 days and nitric acid during this time is either washed out of atmosphere or is removed by surface deposition. Another mechanism associated with lifetime of NO2 is the day-night cycle of its release and capture associated with N2O5. During night, NO2 and nitrate radical (NO3) combine in presence of some catalyst to form N2O5 which acts as a strong reservoir. During daytime, sunlight reverses the process and NO2 is released.

Associated with NOx and its impact on ozone are RO2 reactions which can lead to a wide variety of complex non-methane hydrocarbon reactions. Most well known byproduct of this process is PAN (peroxyacetyl nitrate) which acts as a reservoir for NOx in clean marine air. Its free tropospheric values tend to be in the 25-35 pptv range.

2. Carbon monoxide: In natural atmosphere, CO is created as byproduct of reactions sequence of oxidation of methane during photodissociation of HCHO. There is strong correlation between concentrations of CO and methane in troposphere. Average concentrations of CO are on the order of 30-200 ppbv and its lifetimes are relatively short (about 1-2 months) due mainly to reactions with OH. CO and ozone show positive relationship in areas of higher NOx where ozone is being created photochemically. However, in areas of ozone destruction, where NOx concentrations are less than 0.01 ppbv, CO concentrations are independent of ozone.

3. Methane and Non-methane hydrocarbons (MHC & NMHC): Methane is the most important and most abundant atmospheric hydrocarbon. Its lifetime in troposphere is about 5-10 years. Major sink of methane is its reaction with OH leading to the formation of ozone. Another sink is its gradual transfer to stratosphere through exchange processes across tropopause. Methane then acts as an important factor in strotospheric chemistry.

Non-methane hydrocarbons (NMHC) in atmosphere may also contribute to the formation of ozone. However, which species of NMHCs are important and in what amounts is yet not well established.

Major molecules associated with tropospheric ozone chemistry and their energy requirements are listed in Table- 4.

Table- 4. Energy requirements of some major molecules associated with tropospheric ozone chemistry.

Chemical species        Enthalpy of        Free-energy

                                 formation            of  reaction*

O(3P)                           59.6                  55.4

O(1P)                          104.8

O3                                34.1                   39.0

OH                                  9.3                     8.2

HO2                              ~3.4                    4.4

H2O2                          -32.6               -25.2

H2O                            -57.8               -54.6

N2O                              21.6                20.7

NO2                                7.9                12.3

CH4                            -17.9              -12.1

CO                               -26.4              -32.8

+RO2                             var.             var.

* energy needed to create or destroy chemical bonds. Positive numbers indicate energy must be added to create formation reaction.

+ Complex organic peroxy radicals

Ozone chemistry of stratosphere

Most of the ozone in the atmosphere forms the Ozone layer in the stratosphere at altitudes between 10 and 40 km (100 to 0.1 mb pressure altitude) depending on latitude, just above the tropopause. This layer is crucial for life because only ozone absorbs UV-B radiation between 280-320 nm. UV-A rays between 320 and 400 nm are not affected by ozone while UV-C rays between 200 and 280 nm are absorbed by other atmospheric constituents also beside ozone.

Stratospheric ozone distributions are strongly dependent on stratospheric circulation patterns, varying according to latitude, seasons, short-term meteorological changes and the photochemical processes of formation and destruction. Major driving forces are availability of sunlight and thus of UV radiation and in upper stratosphere (above pressure altitude of 5 mb) the latitudinal temperature gradient which assists ozone transport. The ozone content of stratosphere is highly dynamic and variable. Its concentrations peak around the altitude of 30 km in tropics and around 15 to 20 km in polar regions.

Though hundreds of reactions are known to be involved in the ozone chemistry of stratosphere, only a few can be described properly. The ozone chemistry basically involves two types of reactions: those involved with ozone formation and those involved with ozone destruction. These two types of reactions are important because relationship between stratospheric ozone and climate has been studied particularly in association with ozone depletion and ultra-violet radiation. Another important feature is that above tropopause, liquid water does not play significant role and stratospheric ozone chemistry here is dominated by photochemical reactions.

1. Ozone formation: This itself is a photochemical process involving UV radiation of wavelength less than 242 nm. Though photodissociation of oxygen by UV radiation at less than 175 nm may yield an oxygen atom in excited state i.e. O(1D), such photodissociation is important only in the upper stratosphere because such short wavelength can not penetrate lower into stratosphere.

Thus in upper stratosphere reaction may be:

O2 + hv ( O(3P) + O(1D)

Oxygen atom in excited state on collision with some diatomic molecule (M2) yields oxygen atom in ground state i.e. O(3P):


O(1D) + M2 —————–> O(3P) + M2


while in lower stratosphere reaction is:

O2 + hv (175-242 nm) ———> O(3P) + O(3P)

The oxygen atoms in ground state react with diatomic oxygen molecules to form ozone:

O(3P) + O2 ———> O3

2. Ozone destruction: This involves those reactions which balance the photochemical formation of ozone in stratosphere:

O3 + hv ——> O2 + O(1D)

O3 + O ——-> 2O2

Another additional reaction for removal for oxygen atoms is:

O + O + M ——–> O2 + M

Many analogous reactions involving H, N and Cl radicals also occur in stratosphere:

OH + O3 ———> O2 + HO2

HO2 + O ———> OH + O2

NO + O3 ——–> O2 + NO2

NO2 + O ——–> NO + O2

O3 + Cl ——> O2 + ClO

ClO + O ——> O2 + Cl

All the above pairs of reactions are summed as:

O3 + O —–> 2O2

i.e. each pair of reactions involves destruction of ozone and atomic oxygen while restoring the OH, NO or Cl radical.


Filed under: Atmospheric chemistry — gargpk @ 2:43 pm

Methane is emitted from the earth’s surface mainly due the activity of methanogenic bacteria. Mean rate of its emission is 2 x 1011 cm-2 s-1. It undergoes a complex series of reactions which together constitute the methane cycle in the atmosphere. The cycle may be divided into three main parts:

1. Oxidation of methane and formation of formaldehyde;

2. Oxidation (removal) of formaldehyde and formation of carbon monoxide;

3. Oxidation (removal) of carbon monoxide and formation of carbon dioxide.

Oxidation in various reactions of these three parts of methane cycle in achieved by reaction with OH, O2, or by photochemical oxidation. Reduction at places in the cycle is achieved by reaction with NO and HO2.

Oxidation of methane and formation of formaldehyde

Methane in the atmosphere is first attacked by hydroxyl radical yielding methyl radical and water. Methyl radical, through various oxidation and reduction reactions in which methyl peroxide (CH3O2), methyl hydroperoxy (CH3OOH), methyl oxide (CH3O) are formed, finally yields formaldehyde. In the sequence of these reactions, OH and HO2 radicals used are again formed. The reactions involved in this part of methane cycle are:

1. CH4 + OH ——–> CH3 + H2O –

2. CH3 + O2 + M ——> CH3O2 + M

(M is some molecule acting catalytically and carrying off the excess energy of reaction)

3. CH3O2 + NO ——–> CH3O + NO2

3A. CH3O2 + HO2 ——–> CH3OOH + O2

3B. CH3OOH + hv ———–> CH3O + OH

4. CH30 + O2 ——-> HCHO + HO2

Oxidation of formaldehyde and formation of carbon monoxide

Formaldehyde formed ultimately in the above part of methane cycle is removed by photochemical or chemical oxidation reactions in this second part of methane cycle. A very small part of formaldehyde may be removed from atmosphere through dissolution in rainwater. Ultimate product of chemical removal of formaldehyde is carbon monoxide.

5. HCHO + OH ——–> HCO + H2O

5A. HCHO + hv ——–> HCO + H

6. HCHO + hv ———> CO + H2

7. HCO + O2 ———-> CO + HO2

Oxidation of carbon monoxide and formation of carbon dioxide

Carbon monoxide formed in the second part of methane cycle is finally oxidized by reaction with hydroxyl radical to yield carbon dioxide and hydrogen atom.

8. CO + OH ——–> CO2 + H

The notion of continuity helps in understanding the transfer of material along various reaction pathways in the complex set of reactions of methane cycle given above. Formaldehyde formed by oxidation of methane in atmosphere is removed by four possible processes (reaction numbers 5, 5A, 6 and rainout). The notion of continuity requires that the sum of fluxes through these four pathways is equal to the production rate. From the reactions given above the destruction of formaldehyde can be equated with the production of methane at the surface of Earth or with destruction of methane in the atmosphere. Thus may be written as:

– {d[CH4}/dt} = – {d[HCHO]/dt}

Since washing out with rain (rainout) is very insignificant, equation can be rewritten as:

k1 [OH][CH4] = k5 [HCHO][OH] + J5A [HCHO] + J6 [HCHO]

(subscript numbers refer to reaction numbers given above)

The atmospheric concentration of methane is 1.6 ppm or 4.2 x 1013 cm-3 and of hydroxyl radicals is about 7 x 105 cm-3. Taking the rate constant k1 = 8 x 10-15 cm3 s-1, the destruction rate of methane can be estimated as 2.3 x 105 cm-3 s-1. Furthermore, the above equation can be rearranged as:

[HCHO] = k1 [OH][CH4] / {k5 [OH] + J5A + J6}

This gives the estimate of formaldehyde as 4.3 x 109cm-3 (where k5 = 1.3 x 10-11 cm3 s-1 and J5+J6 = about 4.5 x 10-5 s-1). This estimate of the concentration is a little low but not too far from the typical value of 1010 cm-3 that is observed in the atmosphere.

Notion of continuity can be applied to the formation of carbon monoxide from the oxidation of methane. Carbon monoxide in atmosphere may arise from various sources but the magnitude of natural sources of production of CO can be easily assessed. If small loss of formaldehyde and possibly of methyl hydroxyperoxide (CH3OOH) due to rainout is neglected then CO should be formed at the same rate as methane is released into the atmosphere i.e. at

2 x 10^11 /sq. cm/sec. or about 0.7 x 10^15 g (C) /a. This is larger than the amount which arises from human activities (0.3 x 10^15 g (C)/a).



Filed under: Atmospheric chemistry — gargpk @ 2:16 pm
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Ionosphere is the conducting layer at an altitude of about 80 km and above. This zone of atmosphere was initially probed by radio-waves from ground and later by radio-sounders carried by rockets or direct measurements of gaseous components. Salient features of the chemistry of ionosphere are discussed below.

1. Ionosphere can be differentiated into various layers which represent zones of different electron densities. As a whole, ionosphere is electrically neutral since it also has positive ions like O2+, O+ and NO+. The positive ion chemistry is highly distinctive for various layers of ionosphere.

2. Ionosphere structure shows diurnal and long-term changes. Most important long-term changes correspond to solar sunspot cycle. Changes affect reflection of radio-waves and also alter the concentrations of various species in upper atmosphere.

3. Electrons in ionosphere are produced by photo-ionization. Above the altitude 100 km, this photo-ionization is brought about largely by extreme ultra-violet radiation. At lower altitudes, Lyman-A radiation is important. Some contribution to photo-ionization at somewhat lower altitudes is also made by cosmic rays. However, due to magnetic shielding of Earth, cosmic radiation is only important at fairly high latitudes. Night-time ionization is attributed to a downward flux of protons and radiation from excited species in the upper atmosphere (i.e. UV night glow).

4. In D region of ionosphere, electrons are produced principally by photo-ionization of nitrous oxide because it has lowest ionization potential among dominant species in the atmosphere. However, NO+ is not the most abundant positively charged species in the upper atmosphere. At altitude about 80 km, principal ion is a water cluster or hydrated proton i.e. H+(H2O)2. The charge initially carried by NO+ is transferred to water via an O2+ intermediate.

5. Production of electrons and ions is balanced by loss processes in a quasi-steady-state ionosphere. Loss processes usually involve reduction of photo-electron to thermal energies followed by ion-electron recombination or electron attachment. Typical processes are:

NO+ + e- ——> N + O (dissociative recombination)

O+ + e- ——–> O + hv (radiative recombination)

O2 + O2 + e- ——> O2- + O2 (three-body attachment)

6. In F-layer of ionosphere, positive charge is largely carried by O+ while at lower levels, it is more likely to be present on NO+, O2+ and lower down in atmosphere, on hydrated proton.

7. Though hundreds of reactions are used in descriptions, positive ion chemistry is still poorly understood. D-region of ionosphere is particularly complex because of the presence of an extensive array of negatively charged poly-molecular hydrates of water.

8. E-region of ionosphere is interesting because it sometimes shows thin sporadic layers that appear to be derived from metal-ion chemistry in mid-latitudes. Intensities of these layers show significant increases in response to meteor showers so it is possible that metal ions have extraterrestrial origin. Typical reactions are:

Mg + hv ——-> Mg+ + e-

Mg+ + O2 + M ——-> MgO2+ + M

Mg+ + O3 ——–> MgO+ + O2

The first reaction produces electrons but subsequently they react with charged metal and metal oxide species.

9. In the ionosphere, O+ ions are normally removed through reaction with oxygen and nitrogen:

O+ + O2 ——> O2+ + O

O+ + N2 ——> N2+ + O

But reactions involving hydrogen or water are about 1000 times faster. This leads to considerable reduction in concentration of electrons through following reactions:

O+ + H2O —–> H2O+ + O

O+ + H2 ——> OH+ + H

followed by:

e- + H2O+ —–> H2 + O

e- + H2O+ ——> OH + H

e- + OH+ ——-> O + H

10. Human activities can also affect the ionosphere chemistry. For example, at first launch of Skylab, a large booster operated in upper portion of ionosphere (at altitude 190 km). During the portion of flight through ionosphere, some 1.2 x 1031 molecules of water and hydrogen were released were released due to which electron densities were lowered over a radius of 1000 km around the flight path of the rocket thus creating an electron-hole.


Presence of water and enough light in the clouds may result in the formation of hydroxy and hydroperoxy radicals there. These radicals shall be scavanged by cloud droplets and then could promote a variety of reactions in the droplets.


(a) Oxidation of inorganic species such as ammonia:

NH3(aq) + OH(aq) —— NH2(aq) + H2O

NH2(aq) + O2(aq) —– NH2O2(aq)

NH2O2(aq) + OH(aq) —— HNO2(aq) + H2O

(b) Oxidation of nitrogen oxides:

NO-2(aq) + OH(aq) —- NO2(aq) + OH-(aq)

NO(aq) + OH(aq) —— HNO2(aq)

NO2(aq) + OH(aq) —— HNO3(aq)

(c) Oxidation of sulfur compounds:

H2S(aq) + OH(aq) —— HS(aq) + H2O


Scavenged HO2 radicals have a longer lifetime than OH radicals in the water droplets, so they may be at much higher concentrations there and, therefore, could be important in reactions such as:

(a) Oxidation of SO32-

HO2(aq) + SO2-3(aq) —— SO2-4(aq) + OH(aq)

(b) Generation of Hydrogen peroxide:

HO2(aq) —— H+(aq) + O-2(aq)

O-2(aq) + HO2(aq) —– HO-2(aq) + O2(aq)

H+(aq) + HO-2(aq) —– H2O2(aq)


The possibility of radical chemistry opens up a whole range of organic reaction chemistry also, in particular the oxidation of naturally occurring alcohols and aldehydes e.g.

CH3OH(aq) –oxidant—- HCHO(aq) + H2O

HCHO(aq) + H2O ——- H2C(OH)2(aq)

H2C(OH)2(aq) + OH(aq) ——- HC(OH)2(aq) + H2O

HC(OH)2(aq) + O2(aq) ——- HO2(aq) + HCOOH(aq)

Formic acid, acetic acid and oxalic acid have been detected in the rainwater and point to the possibility of detection of a wide range of dissolved organic substances. These may indicate a complex radical-initiated chemistry that has an important effect on the acidification of rainwater.


Filed under: Atmospheric chemistry — gargpk @ 2:01 pm

Many key reactions in the atmosphere are photochemical reactions which are initiated by absorption of a photon of light. Such reactions can be written as if they were normal chemical reactions by substituting photon (hv) as one of the reactants:

NO2 + hv —— NO + O

The rate constant of such a reaction is given as:

– {d[NO2]/dt} = k” [hv][NO2]

This expression is not very useful because the second-order rate constant k” would probably vary dramatically with the energy of photon involved in the reaction. However, by taking a content flux of photon with respect to wavelength and incorporating it into a pseudo first-order rate constant, the rate expression becomes:

d[NO2]/dt = J[NO2]

J is special first-order constant that embraces the absorption coefficient of the reactant, quantum efficiency of the reaction in question and the solar spectrum and intensity at the altitude and latitude under consideration. Estimates of J for many atmospheric trace gases can be made with a little information on the photochemistry. For example, a typical mid-latitude mid-day value of JNO2 , for the photodissociation of nitrogen dioxide is 5 x 10-3 s-1 which suggests a residence time of 200 s.

Many photochemical are important in the atmosphere as they yield atoms or free radicals and these species are greatly more reactive than the molecular species found in the air. For example, photodissociation of NO2 yields atomic oxygen which can subsequently lead to the formation of ozone:

O + O2 + M ——-> O3 + M

where M is a third body i.e. a molecule such as molecular nitrogen which carries off the excess energy that might disrupt the ozone molecule. The ozone thus produced might further be photodissociated:

O3 + hv ——–> O(3P) or O(1D) + O2

If wavelength of photon is less than 315 nm, the oxygen atom is produced in excited 1D state, otherwise in the 3P ground state. The ground state oxygen will probable recombine with a molecule of oxygen to for ozone again i.e. no net reaction would occur. The excited oxygen atom may be collisionally de-excited to ground state, or more importantly, may react with water molecule providing a source of hydroxyl radical (OH):

O(1D) + H2O ——–> 2OH

Hydroxyl and hydroperoxy radicals in atmosphere

Hydroxyl radical (OH) produced by reaction of excited oxygen atom (formed by photodissociation of atmospheric ozone) with water as described above is probably the most important radical in the chemistry of troposphere. A number of reactions in the troposphere involving hydroxyl radical can produce hydrogen atom or hydroperoxy radical:

OH + CO —–> CO2 + H

OH + O ——> O2 + H

OH + O + M ——> M + HO2

OH + O3 ——> O2 + HO2

H + O2 + M ——> M + HO2

Very quickly a range of radicals and atoms can be generated. These highly reactive species are basic to the gas-phase chemistry of atmosphere. Due to their high reactivity, these species are naturally found in very low concentrations in the atmosphere. Their typical background concentrations are:

Hydroxyl radical – 700000per cubic cm; Hydroperoxy radical – 20000000 per cubic cm

High reactivity of these radicals is indicated by their short residence times. The residence time of OH radical is less than 1 s while that of HO2 radical is perhaps 1 minute.

Reactions of hydroxyl and hydroperoxy radicals with atmospheric trace gases

Reactions of many trace gases found in the atmosphere with the hydroxyl radical exert a profound effect on the composition of atmosphere. Reactions with some of the atmospheric trace gases are discussed below.


Biologically produced sulfur gases are emitted into the atmosphere mainly as sulphides such as dimethyl sulphide, hydrogen sulphide and carbon disulphide. All these react with hydroxyl radical in the atmosphere.

(a) Dimethyl sulphide: This is the major sulphide emission into the atmosphere which reacts with OH radical as follows:

CH3SCH3 + OH —–> CH3SOH + CH3

O3 + CH3SOH —–> CH3SO3H

The product of these reactions is methyl sulphonic acid (CH3SO3H) and most of it persists in the ambient atmosphere though a relatively small amount may be oxidized through sulfur dioxide.

(b) Hydrogen sulphide: This gas in atmosphere is also attacked by OH radical as follows:

H2S + OH ——-> HS + H2O

The resulting bisulphide radical (HS) is oxidized through SO2 in a number of subsequent reactions. The SO2 can also be oxidized by OH and HO2 radicals:

SO2 + OH + M —–> HSO3 + M

SO2 + HO2 —–> SO3 + OH

Bisulfite radical (HSO3) and SO3 react with OH and water respectively to yield sulfuric acid which is the ultimate product of oxidation of atmospheric sulfur.

HSO3 + OH ———> H2SO4

SO3 + H2O ———-> H2SO4

(c) Carbon disulphide (CS2): This has been experimentally shown to be oxidized by OH radical yielding equal proportions of carbonyl sulphide and sulfur dioxide as final products. However, in atmosphere CS2 may not react with OH radical principally and its reactions with oxygen atoms may be more important.


Most of the atmospheric ammonia is removed through dissolution in liquid water in the atmosphere. However, ammonia is also attacked by OH radical though this reaction accounts for only a few percent of the ammonia removed from Earth’s atmosphere:

NH3 + OH ——> NH2 + H2O

Various subsequent reactions are possible:

NH2 + O ——> HNO + H

HNO + O2 —–> NO + HO2

The NO can be oxidized to NO2 which subsequently may react with OH radical to yield HNO3 and this is effectively removed from atmosphere through dissolution in rainwater.

NO2 + OH ——> HNO3


Hydroxyl radical on reaction with carbon monoxide yields carbon dioxide and hydrogen radical.

CO + OH ———-> CO2 + H


Formaldehyde found in trace quantities and formed in various atmospheric reactions is oxidized by OH radical in the following manner:

HCHO + OH ——–> HCO + H2O


Methane is naturally emitted from earth’s surface. In the atmosphere, methane is oxidized by OH radical yielding methyl radical and water:

CH4 + OH ——–> CH3 + H2O

CH3 undergoes following reactions in the methane cycle in the atmosphere yielding CH3O2.

CH3 + O2 + M ——-> CH3O2 + M

CH3O2 reacts with hydroperoxy radical in the following manner:

CH3O2 + HO2 ——–> CH3COOH + O2


In the presence of some suitable molecular species (M), hydroxyl radicals may react with each other to for hydrogen peroxide:

OH + OH + M —— H2O2 + M

Hydroperoxy radical may be a more efficient route for the formation of hydrogen peroxide:

HO2 + HO2 + M —— H2O2 + M

H2O2 is highly water-soluble and a strong oxidizing agent so it probably plays an important role in oxidation processes within water droplets in the atmosphere.

December 23, 2008

Acid rain

Filed under: Acid rain,Air pollution,Atmospheric chemistry,Environment — gargpk @ 3:25 pm

 The natural rainwater moving through atmosphere comes in contact with various chemicals produced and deposited in the atmosphere. These chemicals are produced by various natural (e.g. electrical nitrogen fixation due to electrical lightening), biological (e.g. release of gases in decay and decomposition of organic matter and other biological processes) and geological (e.g. volcanic eruptions and weathering of rocks) processes. Due to this contamination, natural rainwater in perfectly unpolluted areas is also somewhat acidic. The pH of normally clean or ‘pristine’ rainwater is generally agreed by scientists to be 5.6.If rainwater falling in an area has pH value below 5.6, it is called acid rain.

Recent measurements show that rain and snow having pH 4.3 or below fall regularly over many areas of heavily industrialized Northern hemisphere, specially North America, northern and western Europe. Sometimes individual storms under favourable conditions may have may have very low pH values. For example, in 1979, Kane in Pennsylvania, America recorded a rain of pH 2.7 and in same year, , Wheeling in West Virginia had rain of pH 1.5. In Britain, Pitlochry had a rainfall of pH 2.4 in 1974. The acid rains is caused by emission of large quantities of sulphur dioxide and oxides of nitrogen in the atmosphere due to burning of fossil fuel in various industrial and other activities of human beings. Allied to acid rains are phenomena of acid mist and acid fog, both of which come under the category of occult precipitation. The cause of acid mist or acid fog is high concentration of sulphates and nitrates in the form of fine aerosol particles (dust or soot) in wind-driven ground-level clouds which causes condensation of tiny water droplets around these particles. These droplets being tiny fraction of normal rain drops, do not fall as rain water but remain suspended in the atmosphere forming acid mist or acid fog.

The problem of acid rain has attracted worldwide attention only since 1980s. However, the term ‘acid rain’ was first used by first Alkali Inspector of Britain,, Robert Angus in 1872. His work largely remained ignored until 1950s when Canadian ecologist Dr. Eville Gorham undertook detailed studies of rainwater quality and its control in Lake District in north-west England. By mid 1960s, early damage symptoms of acid rains begun to appear in Scandinavia and Swedish worker Svente Oden begain a concerted scientific effort in 1967 to bring awareness about acid rain problem. He is considered to be the father of modern acid rain studies.


Primary pollutants causing acid rain problem are blown over long distances by the wind and thus spreading the problem over whole of the Earth’s surface. However, till now most of pollutants responsible for acid rain problem are produced in the highly industrialized nations, the areas of the impact of acid rains are few, noticeable, few and predictable. Common properties observed in areas affected seriously by acid rain problem are:

  1. Heavy concentration of industries producing pollutants responsible for acid rain problem.

  2. Downwards flow of winds from pollutant-producing areas.

  3. Upland-mountainous position of pollutant-producing areas having thin glaciated bedrock and high rainfall-snowfall.

  4. Numerous lakes and streams and rich forest cover in pollutant-producing area.

Areas sharing the above common properties are termed acid rain hot spots and include many parts of Scandinavia, upland Britain, West Germany and many parts of Northern Europe. Across Atlantic, such areas include Nova Scotia, Canadian Shield around southern Ontario and Quebec, Adriaondack Mountains, Great Smoky Mountains, parts of Wisconsin and Minnesota, Pacific Northwest U.S.A., Colorado Rockies and Pine Barrens of New Jersey. Japanese islands are also included in this category.

In contrast to above areas, there are two types of safe areas where acid rains are not a problem at present. These areas include:

  1. The areas located away from and not downwind of possible source areas and themselves having little polluting industrialization. These areas include almost all of southern hemisphere, tropics and parts of northern hemisphere e.g. northern Russia.

  2. The areas that receive acid rains but have natural resistance to its damaging impact due to buffering capacity provided by the alkaline dust blown from the west. Actually alkaline rains have been reported in Sweden before 1960 in areas with limestone outcrops and cement manufacturing areas. Wind blown alkaline material can de derived from deserts (fine material brought over from Sahara and Gobi deserts has been reported), from wind erosion of top soil alkaline particulate pollutants e.g. soot from smoke-emitting chimneys and agricultural fertilizers.

Geologically, the areas most vulnerable to acid rains fall under three categories:

  1. Glacial areas on granite and other highly siliceous bedrock e.g. quartzite, quartz sandstone, certain gneisses and on materials derived from these.

  2. Areas with thick deposits of siliceous sands e.g. sand plains of Denmark and Netherlands.

  3. Areas with relatively old, highly weathered and leached soils.

Areas having severest acid rain damage are glaciated Pre-Cambrian shield areas of Scandinavia, glaciated parts of upland Britain having thin soils, eastern Canada and resistant Canadian Shield and northwest U.S.A. Problems of acidification develop much acutely on granite and similar other resistant rocks.

Acid rain as global problem

Though at present acid rain problem is mainly concentrated in highly industrialized areas, the long-range transport of concerned air pollutants results in gradual globalization of the problem. As a result of slow transport of acid rain causing pollutants from heavily industrialized areas to areas till now free from this problem, the latter areas are also beginning to show acidification damage. Such damage has been reported from many developing nations like Zambia, South Africa, Malaysia, Venezuela, India and China. Most productive farmlands of China and India, paddy fields of South-east Asia and forests of Amazon in South America have soils which are highly susceptible to acidification.

Global dimension of acid rain problem was established beyond doubt in 1981 with discovery of Arctic haze. It is bluish-gray haze developing in Arctic areas similar to that frequently found over and downwind of large industrial areas in western Europe and eastern North America. Haze layers often cover a horizontal area of upto 1000 km and are caused by scattering of solar radiation by minute suspended particles in the atmosphere. These particles vary in the size range of 0.1-1.0 micrometer and mostly comprise of sulphate aerosols. These aerosols are transported by jet streams in upper atmosphere and may reach upto 8000 km away from their industrial sources. Hazes are found to be thickest in Alaska’s North Slope extending atleast to Norway. Hazes mainly affect visibility and are not as damaging as the smog.


SO2 and oxides of nitrogen (NOx) emitted into the atmosphere due to industrial, commercial and other anthropogenic activities are the basic cause of acid rain formation. Therefore, the problem of acid rains has accompanied the rise of emission of these gases into the atmosphere.

SO2 is emitted from three principal man-made sources:

  1. Combustion of coal produces about 60% of total SO2 emitted into atmosphere.

  2. Combustion of petroleum products which adds 30% of total emission.

  3. Industrial activities like smelting of iron, zinc, nickel, copper ores, manufacture of sulphuric acid and operation of acid concentrators in petroleum industry. These produce the remaining 10% of this gas.

Overall emission of oxides of nitrogen is small in comparison with SO2, their importance in formation of acid rains is very high. Most of the oxides of nitrogen (NO3, NO2, NO etc.) are produced from:

  1. Combustion of fossil fuels.

  2. Industrial chimneys and thermal power stations.

  3. Motor vehicles in urban areas.

Man-made sources of SO2 and NOx emission are point sources (e.g. thermal power stations and industrial chimneys) and the emission from these occurs as a plume of gases. The plume of gases emitted from high stacks usually travels downwind for about 12 km as a straight line without much dispersion. Afterwards, its shape evolves by diffusion and changes progressively downwind into a widening cone. The direction, speed, distance of travel of the plume and its dispersal and diffusion depend upon meteorological conditions such as direction, velocity and pattern of propelling wind, air temperature (especially the vertical temperature gradient), air turbulence and atmospheric stability. Under stable atmospheric conditions, for example, at night over land and during day over snow covered ground, there is very little vertical dispersal for very long distance and the acidification may occur at quite far away place from the source of emission.

Dispersal of the plume of SO2 and NOx occurs in the mixing layer of atmosphere that extends from ground level upto 1-2 km altitude. The dispersal is triggered by diffusion and atmospheric turbulence, normally between 5 to 25 km from the point of source. The rate of diffusion and mixing of oxides into air is faster when flow of air is turbulent. The lower portion of the dispersing cone of oxide plume first touches the ground level at about 5 km distance form the point source while middle and upper portions are thoroughly dispersed in the air leading to dilution and chemical transformation.

The deposition of pollutant oxides from the plume onto the ground is of two types: dry deposition and wet deposition.

  1. Dry deposition: The acidic oxides deposited from the bottom of the plume between 5-25 km from the source in the form of gases and particles constitute the dry deposition. Though such deposition is not acid rain in strict sense, it produces acidification of soils and surface water bodies similar to acid rain. This dry deposition also causes direct SO2 and NOx poisoning of the vegetation. Dry deposition of sulphur and nitrogen oxides and undissolved acids on lakes and steams straightaway dissolve in the water and acidify the water bodies. Such dry deposition on land and on vegetation remains inactive till dew or rainfall when these dry deposited acids dissolve in the dew or rain water and form active acids. Such sudden addition of high concentration of acids into an otherwise stable environment causes acid shocks, acid flushes or acid surges. These terms indicate increasing levels of acidification and decreasing time period in which such acidification takes place. During winters, SO2 and NOx pollutants are dry-deposited on snow and ice in the catchment areas of many lakes and rivers. In the following spring season, when this snow and ice melt, the acids accumulated in the snow and ice over long period are suddenly released over a period of few days to a week causing acid surges in the lakes and streams.

  2. Wet deposition: It is the deposition of acidic oxides of the plume over land or vegetation after being dissolved in the rainwater, snow or ice forming acid rains, acid snow, acid mist or acid fog. Today’s industrial chimneys are normally 100-300 meters high and, therefore, such wet deposition normally occurs beyond 25 km from the point source. The prevailing wind pattern and the length of time over which oxides are transported in the wind system is of great importance in the geographical distribution of acid rains. Longer the SO2 and NOx remain in the atmosphere, greater is the possibility of their transformation to produce sulphuric and nitric acids.

The practice of increasing the height of chimneys and installation of electrostatic precipitators to reduce the air pollution appears to have magnified the problem of acid deposition in two ways. Firstly, tall stacks of pollutant-emitting units now emit pollutant gases at much greater heights so that these gases are now dispersed over much wider areas increasing the geographical extent of acid deposition. Secondly, installation of electrostatic precipitators and other mechanisms to remove alkaline particulates in chimneys has resulted in increased emission of acidic gases. It is because prior to installation of such mechanisms, acidic gases were neutralized to a large extent by alkaline particulates being emitted alongwith them.


Strictly speaking acid rain is a term which indicates a wide variety of mixtures of acids and oxides in the rainwater. For example, rainwater of pH 4.5 may contain a high sulphur content, high nitrogen content or any combination of the two. Acidity of rainwater results from chemical transformations of a large number of acidic ions added to the atmosphere from natural sources (e.g. sea salts, volcanic emissions, biogenic emissions, soil etc.) and by human. Major such ions can be categorized as following:

  1. Inorganic ions: These include trace metal ions which often act as catalysts to quicken the acidity processes. At coastal sites, corrections for the impact of seawater on rainwater quality have to made before accurate assessment of the role of land-based sources can be made. In individual locations, rainwater quality may be strongly influenced by local sources.

  2. Organic ions: These are important alongwith local biogenic sources in affecting the precipitation quality, particularly in tropics.

Table 1. Major inorganic and organic ions and molecules affecting rainwater acidity.

Ion or molecule

General source




H+ (cation)


SO42- (anion)


NO3 (anion)


Cl (anion)

NH4+ (cation)


Ca2+ (cation)


K+ (cation)


Mg2+ (cation)


Na+ (cation)




H2O2 (molecule)

O3 (molecule)

Fe3+ (cation)

Mn2+ (cation)

NO2 (molecule)



Aqueous chemistry


Combustion of fossil fuels, ocean and soil processes

Agriculture, fossil-fuel burning


Ocean, some industries

Agriculture, decay processes, industry

Soil, agriculture


Soil, agriculture


Ocean, soil, agriculture


Ocean, industry




Aqueous chemistry

Atmospheric chemistry

Soil, industry


Fossil-fuel burning



Amount directly proportional to rainwater acidity

Strong acid, gas and liquid reactions

Strong acid, gas and liquid reactions

Acid, mainly gas reaction

Neutralization of anions


With carbonates, act as buffer acidity

With carbonates, act as buffer acidity

With carbonates, act as buffer acidity

With carbonates, act as buffer acidity



Major at all rainwater pH







HCOOH (molecule)

CH3COOH (molecule)





Weak acid

Weak acid


The steps involved in each chemical process contributing to rainwater acidity depict a multitude of pathways with many of the steps being reversible and many of the steps exhibiting highly complex chemistry. Thus the overall chemistry of acid rain is extremely complicated because of the very large number of chemical interactions involved. Moreover, exact chemical composition of acid rain is not same in every area. It varies from place to place depending upon the proportion of different oxides present and the chemical transformations they have undergone during their stay in the atmosphere. Although a variety of natural and man-made oxides contribute to rainwater acidity through variety of chemical pathways, most important pathways are those associated with two major acidic gases i.e. SO2 and NO2 added to atmosphere from various polluting sources. The complex pattern of acid deposition has following six stages:

  1. The atmosphere receives SO2 and NOx from natural and man-made sources.

  2. Some of these oxides fall on the ground as dry deposition within 5-25 km from their parent sources.

  3. Formation of photo-oxidants like ozone, is stimulated in the atmosphere.

  4. The photo-oxidants interact with SO2 and NOx to produce acids (H2SO4 and HNO3) by oxidation.

  5. The oxides of sulphur and nitrogen, photo-oxidants and other gases (including NH3) dissolve in the cloud and rain-droplets to produce acids (H+ and NH4+) and sulphates (SO42-) and nitrates (NO3).

  6. Acid rain containing ions of sulphate, nitrate ammonium and hydrogen falls as wet deposition.

The most important step in this chain of reactions is the catalytic conversion of SO2 and NOx. This may take from a few hours to a few days in the atmosphere and can not occur without photo-oxidants (precurssors). Ozone is the most readily available and abundant photo-oxidant in the atmosphere . Hydrocarbons and NO added to the atmosphere as pollutants are the two main precurssors of ozone. The acid rain is the final product of the loading of SO2 and NOx coupled with photochemistry and physical dynamics of stratosphere.

Acid gases like SO2 and NOx are transformed into dilute acids in the rainwater by following three major types of reactions:

  1. Homogeneous gas-phase reactions: These reactions occur in the dry atmosphere associated with photolytic oxidation processes.

  2. Homogeneous aqueous-phase reactions: These occur between individual species in a liquid medium such as cloud or raindrop.

  3. Heterogeneous aqueous-phase reactions: These occur during adsorption of acid gases on solid surfaces and are extremely complex. These reactions probably assist in creating rainwater acidity but are not considered to be as important as other two types of reactions in the overall chemistry of acid rains.

The relative importance of any chemical process operating in the atmosphere depends strongly on the meteorological conditions such as the presence of clouds, relative humidity, intensity of solar radiation, temperature etc. Following two factors are crucial to the operation of each process:

  1. Time available to complete secondary chemical reactions.

  2. Availability of excited ions and catalysts to assist the reactions.

Homogeneous gas-phase chemistry

In dry atmosphere, most of the acid gas reactions leading to formation of acid ions such as sulphates and nitrates involve excited molecules, atoms, free radicals and sunlight. The OH radical is particularly important in such reactions. Following main such chemical pathways lead to eventual formation of sulphuric and nitric acids in rainwater:


Very slow reaction:

2SO2 + O2 ——- 2SO3

Unstable compounds:

OH + SO2 + M —– HOSO2 + M

HOSO2 + O2 —- HO2 +SO3

(M = catalyst; often Fe3+ or Mn2+)

Very fast reaction:

SO3 + H2O — H2SO4


Very slow reactions: ( ppb concentrations are reached in many days)

2NO + O2 2NO2


HO2 + NO —– OH + NO2

2NO2 + H20 —— HNO3 + HONO

Factors affecting homogeneous gas-phase reactions

  1. Interfering substances: Oxidation of SO2 and NO2 in the atmosphere is relatively a slow process and there may be several substances causing interference along the way. For example, HOSO2, which is a very unstable substance, may react with CO, NO, water vapour, various hydrocarbons and other chemical species and block the reaction described above.

  2. Catalysts: The reactions between SO2 and NO2 with O2 in the dry atmosphere are considered to be so slow without catalysts that the eventual output of acid is very small. Reactions with the addition of catalysts and free radicals are the main sources of ions leading to acidity of rainwater.

  3. OH radical: Oxidation rates of SO2 and NOx in a cloud-free atmosphere are highly variable and strongly dependent on the concentration of OH radical. If concentration of OH radical is relatively high (on the order of 9×106 mol cm-3), oxidation of SO2 to SO42- is approximately 3.7 +/ 1.9 % per hour. Conversion of NO2 to HONO2 is much more rapid reaction; its rate being about 34 +/ 17% per hour. With lower OH concentrations, the conversion rate is reduced and SO2 converts at a rate of about 0.7% per hour or about 16.4% per day. NOx conversion is at much faster rate and the rates vary between 6.2% per hour and 100% per day. In winters, conversion rates are 0.12% and 1.1% per hour respectively. At night, when OH concentrations are at minimum, conversion rates are sharply reduced.

Homogeneous aqueous-phase reactions

The species of sulphur and nitrogen can be incorporated in liquid water droplets in several ways e.g. (I) they may have high solubility in water; (ii) they may attach through diffusional processes; (iii) they may be incorporated through impactations and collisions and (iv) acid aerosol species may act as nuclei for formation of water droplets. Most important aqueous –phase reactions in acid-rain chemistry are as following:


SO2 + H2O <—- SO2.H2O

SO2.H2O H+ + HSO3

HSO3——— H+ + SO3

O2 + 2HSO3 ——- 2H+ + 2SO42- (Reaction slow without catalyst)

H2O2 + HSO3 ——– H+ + SO42- + H2O (Reaction is rapid)

O3 + HSO3 —— H+ + SO42- + O2 (Reaction is rapid if pH>4.5)


NO2 + O3 —– O2 + NO3

NO3 + NO2 + M <—— N2O5 + M (M = Catalyst; often Fe3+ or Mn2+)

N2O5 + H2O ——- 2H+ + NO3 + NO2

Factors affecting homogeneous aqueous-phase reactions

  1. Reaction medium: Conversion of acid ions is much faster when reaction medium is water. At the droplet scale, sequence of conversions might be:

    1. Initial diffusion of gas to the droplet interface.

    2. Transfer across the interface into the droplet.

    3. Swift aqueous-phase equilibrium.

    4. Aqueous-phase reactions and concurrent diffusion.

  2. Catalysts: In liquid water, catalysts are very important in determining the speed of conversion process. Models using proper chemical conversion estimates indicate that, with the exception of H2O2, impact of other catalysts is highly dependent on the pH level in the water. If pH of the droplet is of the order of 5.0, then conversion rates are significantly increased in the presence of O3, Fe3+, Mn2+ and other ions. However, at pH level of 4.5, trace metal ions contribute only about 1% per hour to the conversion process and the impact of ozone drops to about 10% per hour. At pH level of 4.0, trace metal ions have negligible impact and ozone adds only about 1% per hour to conversion process. This occurs because, in part, solubility of SO2 in water decreases with increasing H+ concentration.

  3. Hydrogen peroxide: It enhances the rate of conversion of SO2 to SO42- independently of the pH level in water droplets. H2O2 dominates the aqueous chemistry process and may increase the conversion rates to 100% per hour depending upon the cloud type, altitude and other meteorological conditions until it is fully exhausted. Afterwards, ozone becomes the dominant catalyst of conversion reactions. H2O2 is not important in the formation of NO3. Favourable conditions for the formation of H2O2 are low NOx concentrations and high concentrations of hydrocarbons and aldehydes in the atmosphere. The conditions favourable for ozone formation are unfavourable for H2O2 formation.

  4. Cations in solution: The rates of formation of SO42- and NO3 may be altered by cations in solutions, particularly by ammonium (NH4+). The cations may increase the rate of oxidation of SO2 by more than an order of magnitude. Ammonia can dissolve as a gas in water droplets and thus directly reduce the rainwater acidity. Presence of extra cations enhances the impact of catalysts, especially at pH above 4.5. This results in formation of disproportionately high amounts of SO42- and NO3 in presence of cations in solutions than in presence of free H+. Ammonium seems to increase the formation of SO42- most in spring when concentrations of both NH4+ and H+ are highest. It has been suggested that about 50% increase in NH4+ in Europe since 1950s may have had some impact on the change in SO42- in rainwater relative to H+. If soil dust rich in cations like Ca2+ and Mg2+ is loaded into the atmosphere, these cations neutralize the strong acids and the rainwater tends towards alkalinity. For example, in India, strong acidic ions in atmospheres around urban areas are heavily neutralized by such soil dust.

  5. Formation of NO3: In areas where concentrations of hydrogen peroxide and ozone are negligible, formation of NO3 can control the production of sulphuric acid in atmosphere. The H2O2 is not important in the formation of NO3. Though very little is known about conversion of NOx in aqueous environment, N2O5 is supposed to play important role and perhaps of NO3 is directly formed from it depending on the relative concentrations of NOx and NO3. In the night, reaction of oxides of nitrogen with ozone can produce significant amounts of NO3 because of the absence of its photochemical destruction.

  6. Season and time of day: Season and time of the day have important impact on acidity of rainwater and cloud-water due to following important reasons:

  1. Difference in pollutants and ions: There are generally different mixtures of pollutants and ions available for acid conversion at different times of day and in different seasons.

  2. Difference in gas-phase reaction rates: In winter, available solar energy is weaker and, therefore, gas-phase chemical reactions are slower than in summer.

  3. Difference in concentrations of catalysts: In winters, oxidation in clouds generally decreases because concentrations of appropriate catalysts are lower than in summers. For example, levels of H2O2 may be about 16 times higher in summers (about 4.8 ppbv) than in winters (about 0.3 ppbv). This high H2O2 concentration in summers enhances the formation of SO42- in that season. At night, conversion of SO2 to SO42- may reach 10% per hour in good catalytic conditions such as low stratus clouds over water.

  4. Difference in photochemistry: The photochemical production and destruction of chemical species in atmosphere depends on the availability and intensity of solar radiation and, therefore, may affect their concentrations during day and night. For example, concentration of NO3 increases considerably at night when it is not being destroyed photochemical reactions.

  5. Types of clouds and precipitation: Mechanism of removal from the clouds may vary by the types of clouds and precipitation. In the clean background air of southern hemisphere, gas-phase and aqueous-phase reactions are almost equal in importance. However, in northern hemisphere, particularly in winter season, aqueous-phase reactions become dominant.

Chemistry of acid fog

More recently, measurements at sites in parts of Europe, California and eastern U.S.A. have shown that in most circumstances, acid fog and water in low clouds has a lower pH value than equivalent acidic rainwater. Average pH values of acid fog in areas of heavy air pollution are about 3.4 and range from 2.8 to over 5.0 On average, mean concentrations of H+ and acid ions are 3 to 7 times higher in fog-water than in equivalent rainwater. Acid fog-water also has higher concentration of anions and cations. There are following five main reasons for the above describe differences:

Fog being located nearer to the ground, is often exposed to higher pollutant concentrations for longer periods of time than the rainwater during below-cloud scavenging. This exposure allows more time for extensive aqueous-phase chemical processes to take place.

Smaller fog and mist particles saturate with gaseous pollutants more quickly than the larger raindrops, allowing greater aqueous ion production.

The smaller droplets in fog and mist have a greater combined surface area compared to raindrops. As a result, acid gas diffusion is enhanced and higher concentrations of the resultant ions are produced.

The fog remains in the air mass in which it is formed while precipitation is often associated with changing air masses in frontal situations when much of the gaseous and aerosol material in the atmosphere is removed.

Pollutant aerosols originating several hundred kilometers away often act as nuclei for fog or cloud droplets and enhance aqueous chemical processes. The size and number of water droplets formed and the resultant chemistry depend on the number of aerosol nuclei available in the cloud. Greater number of these generally produce smaller and more numerous fog droplets. Ion concentrations in mist tend to be lower than in fog because mist contains a lesser number of droplets and this limits the chemical reactions.

However, fog-water shows wide variations in ion concentrations between sites and events. In stable atmosphere, low altitude fog masses are more likely to interact with pollutant emissions near the surface e.g. NOx from automobiles. On the other hand, mountain fogs occurring in a well mixed atmosphere and at times, isolated from low-altitude pollutant emissions due to inversions, tend to be cleaner having pH values of 5.0 and above. On minor scale, dew from polluted atmosphere can also be acidic with free H+ comprising about 80% of acidity while species of sulphur and nitrogen may contribute about 60% and 30% respectively to the acidity. 

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