Environment of Earth

March 11, 2008


Filed under: Atmospheric chemistry,Environment — gargpk @ 2:24 pm
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The atmosphere is made up of a large number of gaseous constituents and in the atmosphere, a large variety of chemical reactions are constantly going on amongst its constituents. The products of one reaction are reactants for other reactions i.e. chemical reaction constitute many of the sources and sinks of gases in the atmosphere. Since the study of atmospheric chemistry mainly attempts to understand the chemical kinetics of atmosphere, a study of the gas-phase reaction rates becomes very important.

Gas-phase reaction rates

For understanding gas-phase reaction rates, following simple reaction may be considered:

2NO + O2 ——-> 2NO2

The reaction rate (R) is given by:

R = – {d[NO]/dt} = – {d[O2]/2dt} = d[NO2]/dt

In atmospheric context, concentrations are usually expressed as number of molecules of gas per cubic centimeter (cm-3). Therefore, the rate of above reaction may be expressed in the form:

R = k[O2]1[NO]2

Where, k = reaction rate constant; exponents denote the order of reaction.

This reaction is second-order with respect to nitric oxide, first-order with respect to oxygen and third-order overall i.e. sum of individual orders. As the units of reaction rate are given in concentration per unit time (cm-3 s-1 ), the rate constant for the above reaction will have the units of cm-6 s-1.

The consideration of reactions of various orders shows that:

(a) Rate constant for first-order reaction will have the units of s-1.

R = – {d[A]/dt} = k[A]t

or, d[A]/[A]t = – k dt

On integration, it gives:

ln[A]t = – kt + c

c is concentration of A at t = 0 and is expressed as [A]0. Thus:

ln{[A]t/[A]0} = – kt

or, [A]t = [A]0 e-kt

Similar expression for reactions of other orders may be given:

Order Differential form Concentration relationship Units

0 – d[A]/dt = k [A]t = [A]0 – kt cm-3 s-1

0.5 – d[A]/dt = k[A]0.5 [A]t = {[A]00.5 – kt/2}2 cm1.5 s-1

2 – d[A]/dt = k[A]2 [A]t = [A]0/{kt[A]0 + 1} cm3 s-1

(b) In case of second-order reaction, if concentration of one reactant is significantly higher than that of other reactant then the reaction can generally be treated as a first-order reaction with respect to reactant at low concentration. For example, reaction of ozone with nitric acid in atmosphere is a second-order reaction:

NO + O3 —–> NO2 + O2

and – {d[NO]/dt} = k”[NO][O3]

k” is the second-order rate constant and for this reaction, its value is 1.8 x 10-14 cm3 s-1 at 300 K. Concentrations of NO and O3 may be assumed as being 0.1 ppb and 15 ppb respectively that are the typical concentrations for lower atmosphere. These concentration units have to be made compatible with those of rate constant. At atmospheric pressure of 1.0, one cm3 of gas contains 2.7 x 1019 (Loachmidt number). This number has to be multiplied by partial pressure of the gas to give appropriate units. Thus, the concentrations of NO and O3 come to be 2.7 x 109 cm-3 and 3.9 x 1011 cm-3 respectively. Using these concentrations, the rate of reaction comes to be 1.8 x 107 cm-3 s-1. This rate is quite high compared with concentration of NO and so the concentration of NO will decline quite rapidly in a closed system. On the other hand, concentration of ozone is very much greater than that of NO and will remain relatively constant. Therefore, the concentration of ozone may be incorporated as a constant within the rate constant and the rate expression can be given as:

– {d[NO]/dt} = k’[NO]

where, k’ is first-order constant given by k”[O3]. The value of this pseudo-first-order rate constant is 0.007 s-1 at the concentration of ozone being considered.

Apparent reduction in the reaction order of a system may occur in other ways also. All that is required for it is that the concentration of one reactant should remain constant. If a reactant is being catalyst or continuously replaced in the system, this might be the condition.

(C) When reactions can be reduced to first-order systems, their study becomes convenient. In such systems, concentration of the reactant will halve over a constant time period regardless of its initial concentration. This period of time is termed half-life (t0.5) and is related to the first-order constant by the expression:

t0.5 = ln(2)/k’ = 0.693/k’

In second-order systems, half-life is dependent on the second-order reactant and relationship between t0.5 and second-order rate constant is given by the expression:

t0.5 = 1/k [A]0

For reaction orders greater than unity, higher the concentration, shorter is the half-life.

Atmosphere as steady-state system

From the above short discussion of chemical kinetics, great stability of the atmosphere would not be expected since it is not in a state of chemical equilibrium. However, the natural atmosphere appears to be quite stable. This apparent stability of atmosphere is because it is in steady state in which various chemical species are continuously being added and removed. The steady-state situation of the atmosphere is maintained by relative constancy of the input and output i.e. by the fact that the rates of the addition/production and removal/destruction of chemical species are equal.

For describing steady-state systems, the residence time or mean lifetime of a chemical species is a useful parameter. Residence time () is easily obtained from the first-order rate constants.

The flux of a material from a system is given by:

Fo = A/

Considering the flux as material lost in a given volume through a chemical reaction of first-order and writing the amount as concentration:

Fo = – {d[A]/dt}


d[A]/[A]t = -kdt


d[A]/dt = -k[A]t


Fo = [A]/ = k[A]


 = 1/k’

Concentration [A] is assumed to remain constant due to continual input of material. This is true of a steady-state system. Thus, the residence times are useful for describing steady-state situations. With steady-state assumption for the oxidation of nitric oxide by ozone discussed above, the calculated value of pseudo-first-order rate constant is 0.007 s-1. This value indicates that nitric oxide has a residence time of about 150 s in the atmosphere.

Various chemical species are continuously added into the atmosphere from the lithosphere or biosphere and are destroyed through complex reaction systems in the atmosphere. Thus atmosphere is a steady-state system and quite complex reaction systems in the atmosphere (e.g. Methane cycle discussed later) can be easily understood if they are assumed to be in the steady state.


Water is present in the atmosphere as suspended droplets and atmospheric gases are dissolved in this atmospheric water. The dissolution of atmospheric trace gases into suspended droplets is one of the most important controls on rainfall chemistry.

Henry’s Law describes the solubility of gases in water and states that at equilibrium the partial pressure of a gas above a solution of the gas is proportional to the concentration of the gas in the solution. However, in much of the atmospheric chemistry, it is useful to imagine the relationship between the gaseous and liquid phase concentrations in terms of a equilibrium of the type:

A(g) = A(aq)

Where, A(g) and A(aq) represent the concentrations of substance A in gaseous and aqueous phases respectively.

By writing the Henry’s Law constant (KH) as the equilibrium constant for this reaction and using pressure to describe the concentration of A in the gaseous phase:

KH = [A(aq)]/pA

If units of pressure and concentration are taken as atm and mol l-1 respectively, Henry’s Law constant will have the units of mol l-1 atm-1. It is clear that larger the value of KH, more soluble the gas will be. Therefore, H2O2 is highly soluble and its large amounts can dissolve in the clouds and rainwater droplets. KH values for some atmospheric trace gases are given in the Table-1.

Table-1. KH values for some atmospheric trace gases at 288 K


Gas                                       KH (mol/l/ atm)


Hydrogen peroxide            2 x 10^5

Dieldrin                                   5800

Lindane                                   2230

Ammonia                                     90

Aldrin                                            85

DDT                                                 28

Sulfur dioxide                              5.4

Formaldehyde                             1.7

Mercury                                            0.093

Carbon dioxide                              0.045

Acetylene                                        0.05

Nitrous oxide                                 0.034

Ozone                                                0.02

Nitric acid                                      0.0023

Methane                                           0.0017

Oxygen                                             0.0015

Nitrogen                                          0.001

Carbon monoxide                       0.001


Henry’s Law constant accounts only for simple dissolution of gases and not for the condition where there is hydrolysis after dissolution. For example, formaldehyde dissolves in water and subsequently hydrolysis to methylene glycol according to the following equations:

HCHO(g) ===== HCHO(aq)

HCHO(aq) + H2O ====== H2C(OH)2(aq)

Thus, the apparent solubility of formaldehyde in water is greater than that expected from Henry’s Law constant. The total amount of formaldehyde dissolved (T(HCHO) in solution will be:

T(HCHO) = [HCHO(aq) + [H2C(OH)2(aq)]

The concentration of methylene glycol will be related to the aqueous formaldehyde by the Laws of mass action:

K = [H2C(OH)2(aq)]/[HCHO(aq)]

Where, K is the equilibrium constant fro hydrolysis reaction. This gives:

T(HCHO) = [HCHO(aq)]+ K[HCHO(aq)]

Since [HCHO(aq)] is known from Henry’s Law, above equation may be written as:


In case of formaldehyde, K is about 2000, i.e. gas is readily hydrolyzed by water so most of it will be found in aqueous solution as methylene glycol rather than as formaldehyde. This makes formaldehyde rather more soluble. The KH is about 1.7 mol l-1 atm-1. At equilibrium with atmosphere pHCHO at 10-9 atm, total concentration of the formaldehyde derived carbon would be predicted to be about 3.4 x 10-6 mol l-1.

Dissolution and hydrolysis of formaldehyde is a rather simple case. Many other atmospheric gases such as carbon dioxide, sulfur dioxide and ammonia undergo more complex hydration reactions and the pH of rainwater is significantly influenced by sets of these hydration reactions.

(I) CO2 + H20 ==== H2CO3(aq)

H2CO3(aq) ==== H+(aq) + HCO3-(aq)

HCO3-(aq) ==== H+(aq) + CO32- (aq)

(ii) SO2(g) + H2O ==== H2SO3(aq)

H2SO3(aq) ==== H+(aq) + HSO3-(aq)

HSO3(aq) ==== H+(aq) + SO32- (aq)

(iii) NH3 +H2O ==== NH4OH(aq)

NH4OH(aq) ==== NH4+(aq) + OH-(aq)

Henry’s Law constant and equilibrium constants for these reactions i.e. KH, K’ and K” respectively are given in the following Table- 2.

Table-2. KH, K’ and K” values for some important atmospheric gases.


Gas KH (mol/l/atm) K’ (mol/l) K” (mol/l)


Carbon dioxide 0.045 3.8 x 10^7 3.7 x 10^11

Sulfur dioxide 5.4 2.7x 10^2 2.7 x 10^7

Ammonia 90 1.6 x 10-5 –

Though sulfur dioxide is a soluble gas, it does not all dissolve in liquid phase of a system of cloud-droplets suspended in air. The ratio of the volume of water and the volume of air is quite low usually being less than 10-6. Therefore, most of the mass of sulfur dioxide remains in gaseous phase in a cloud. Among common atmospheric trace gases, hydrogen peroxide is probably soluble enough to partition predominantly into liquid phase. Under the acidic conditions, ammonia may also be found predominantly in the liquid phase.

Transfer of gases to liquids

When a gas is in high concentration in the atmosphere and at very low concentration in water droplet, there occurs a flux of the gas to the water. The flux of gases across an air-liquid boundary is usually described in terms of a two-film model. This model assumes that there are thin boundary layers on either side of the gas-liquid interface and transfer through these layers is governed by diffusion. As diffusion is a slow process compared with turbulent transport, transfer through the still boundary layer limits the flux of gases to water bodies. In gas-liquid transfer, the total resistance to transport will be the sum of the individual resistance of the two layers. For most of the important atmospheric gases (except perhaps formaldehyde), one of these two resistance is greater than the other resistance. Therefore, the transport of gases across gas-liquid interface may be of two types: gas phase controlled transport and liquid phase controlled transport.

(I) Gas-phase controlled transport: Such transport generally occurs in case of highly soluble gases. The flux (F) of a highly soluble gas across a boundary layer at the air-water interface is given by:

F = k c

where, k = exchange constant having units of ms-1 and c = difference in concentration of gas between the bulk gaseous concentration and the gaseous concentration at the liquid surface. The exchange constant is the reciprocal of the resistance ® and can be obtained from the diffusion coefficient of the gas (D) in the boundary layer (z) i.e.

k = 1/r = D/z

As resistance can be summed, exchange constants must be summed as reciprocals i.e. 1/K = 1/k1 + 1/k2 + …., where K = exchange constant of whole system of boundary layers.

(ii) Liquid-phase controlled transport: Such transport occurs mainly in case of less soluble gases. In such case, c represents the difference between concentration in the liquid at the surface and the concentration in the bulk of liquid. Transfer into liquids can be fast in another way also. Rapid reaction of gas in the liquid phase lowers the liquid-phase boundary layer and transfer becomes gas-phase controlled. Sulfur dioxide is such a gas that is rapidly hydrolyzed to bisulfite and sulfite ions in water so that its dissolution in the water is under gas-phase control and is quite rapid (about 0.5 to 1.0 cm/s). This hydrolysis is only significant for sulfur dioxide at pH>3.0 and so in very acidic solutions, the dissolution will remain under liquid-phase control.


The dissolution of gases in water droplets suspended in the air can be considered as a two step process:

(a) Transfer of gas from bulk atmosphere to the surface of droplet

(b) Mixing of gas within the droplet

Gas-phase transport processes are usually fast so the rate of transfer is limited by the mixing within the droplet. If gas is quite rapidly transferred through the gas-phase boundary layer then the surface of droplet can reach equilibrium with the gaseous phase. More gas can dissolve in the droplet after dissolved gas is mixed inward from the surface of droplet towards its center. In the droplet falling through atmosphere, mixing may also occur through convective stirring. If droplet is stagnant, mixing occurs only through the slower diffusive processes. Diffusive transport within a sphere is given by:

Mt/M = 1 – (6/)  1/n2 exp(-Dn22t2/r2)


where Mt and M are masses of the substance at time t and at equilibrium respectively; n is an integer, r is radius of the sphere.

Typical value for the radius of suspended droplets in air is about 50m while value of diffusion coefficient typical for many dissolved gases is about 10-9 m2s-1. These values suggest that 50% saturation of droplet is achieved in 0.3 second indicating that equilibration of water droplets with atmospheric gases is quite rapid. Large suspended drops may take longer to equilibrate but their equilibration can be quite rapid if they are falling and are being stirred by the air flow on their surface.


The dissolution of soluble atmospheric trace gases into droplet results in much increased concentration of those gases in the small volume of water droplet. This increased concentration of gases allows many new opportunities for chemical reactions inside the droplet despite the fact that only very soluble gases are found partitioned predominantly into the droplet phase. Important such reactions are discussed below.

1. Oxidation of sulfur dioxide: This has been the most frequently studied reaction in aqueous atmospheric droplets due to its importance in acid rain problem. Oxidation of sulfur dioxide by oxygen is very slow in absence of catalysts. Only through the presence of catalysts, such as iron or manganese, such oxidation can be fast enough to be important in atmospheric droplets. At typical acidity (i.e. about pH 5.0) of atmospheric aerosols, sulfur dioxide will be present mainly as bisulfate ion (HSO3-) so the oxidation reaction may be given as:

HSO3-(aq) + 0.5O2(aq) –Fe, Mn———–> SO42- + H+(aq)

Despite many attempts, it is still not clear whether iron or manganese present in rainwater is principally responsible for oxidation of sulfur dioxide.

In remote areas these catalysts, though abundant from crustal sources, may not be in a form that is soluble enough to promote the oxidation reaction. In such areas, H2O2 and O3 may be the oxidants though they are present in atmosphere at quite low concentrations:

HSO3-(aq) + H2O2(aq) ——–> SO42- (aq) + H2O + H+(aq)


HSO3-(aq) + O3 ——–> SO42- (aq) + O2 + H+(aq)

The product of these reactions is sulfuric acid that being much stronger acid than sulfurous acid is responsible for appearance of proton on the right side of above reactions. As oxidation proceeds, droplet becomes more acidic than it was owing to sulfurous acid alone: first due to production of sulfuric acid and secondly from dissolution of more sulfur dioxide to replace the oxidized gas. The solubility of sulfur dioxide is lowered with the increase in acidity but in case of catalyzed reaction, the actual rate of oxidation slows down too. Thus, the oxidation reaction can rapidly come to a standstill. However, many workers consider H2O2 to be a very effective oxidizing agent of SO2 in atmosphere because the rate of oxidation by it actually increases under acidic conditions. So the oxidation rate is not slowed down as the reaction proceeds and more H2SO4 is produced.

2. Oxidation of nitrogen oxides: Nitrogen oxides may be oxidized in droplets to a lesser extent to form nitric acid. This system has not been studied in much detail. There is also the possibility of the dissolution and reduction of two nitrogen oxides:

NO2(aq) + NO3(aq) + H2O ———-> 2H+(aq) + 2No3- (aq)

3. Reaction with chloride ion: The production of sulfuric or nitric acid may result in an important subsequent reaction if chloride ion is present in high concentration in the atmosphere. Hydrogen chloride is more volatile than other strong acids found in aerosol droplets, so it may be lost from the droplet according to the reaction:

H+(aq) + NaCl(s) ———> HCl(g) + Na(aq)

Study of the Na:Cl ratio in maritime aerosols has provided evidence of the occurrence of this reaction. However, quantitative description of the reaction is rather difficult because it probably occurs in droplets which have nearly evaporated to dryness. These aerosols will have a very high salt concentration. Under such situation, behaviour of sodium departs from ideal condition meaning that thermodynamic predictions made using equilibrium constants obtained from low concentrations may be wrong i.e. the solubility of gases in saline droplets may not follow Henry’s Law.

4. Dissolution and hydrolysis of carbonyl sulphide: This reaction can be important in remote areas in generating sulfuric acid. Initial Hydrolysis step would be:

OCS(aq) + H2O ———> CO2(aq) + H2S(aq)

This would be followed by the oxidation of hydrogen sulphide through to sulfuric acid.

5. Reactions of photochemically generated species: Hydroxy and hydroperoxy radicals are produced photochemically in the atmosphere and these radicals are scavenged by the cloud droplets. These radicals can then promote various important reactions in the droplets. Such reactions have been discussed later.


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 Table- 3.

Table- 3. 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.


In case of carbon dioxide and sulfur dioxide gases, the K” is very much smaller than K’ and may be neglected at acidic pH values. Thus the pH of a droplet of water in equilibrium with atmospheric carbon dioxide can be determined by combining two equilibrium constant equations, one governing the dissolution and the other the first step in the dissociation. The concentration of HCO3- will be:

[HCO3-] = KH K’ pCO2 / [H+]

If dissociation of carbon dioxide is the only source of hydrogen ions in the system, then [HCO3-] = [H+] so that:

[HCO3-] = [H+] = (KHK’ pCO2)0.5

By substituting appropriate values of equilibrium constants and use of a carbon dioxide partial pressure of 3.4 x 10-4 atm, a hydrogen ion concentration of 2.3 x 10-6 mol l-1 or a pH of 5.6 will be obtained. In remote regions, pH of pure rainwater may be close to this value and been assumed to be the pH of normal rainwater. However, trace amounts of other compounds can affect the rainwater pH. For example, sulfur dioxide concentration of having partial pressure of 5 x 10-9 atm in air will give an equation analogous to the one used for CO2:

[HSO3-] = (KH K’ pSO2)0.5

Use of appropriate values will give a rainwater pH of 4.6. Thus even at low concentrations sulfur dioxide has profound effect on pH despite carbon dioxide being present in much higher concentration. High solubility of sulfur dioxide gas and high dissociation constants for it make it more effective at acidifying water droplets than carbon dioxide. Its oxidation to sulfuric acid is comparatively easy and yields a further proton. Therefore, this gas has even more dramatic effect on rainwater pH.

Equations describing the dissolution of an acidic gas such as sulfur dioxide show that presence of acids already in solution will depress the solubility of the acidic gas while alkalis will enhance it. Ammonia is common alkaline gas in atmosphere and it will neutralize the dissolved acids, particularly the sulfuric acid. This means that ammonium salts, particularly ammonium sulphate, that are present in the atmosphere also effect the rainwater pH. In a system of water in equilibrium with CO2, SO2 and NH3 gases at the same time with pCO2 = 3.4 x 10-4 atm, pSO2 = 5 x 10-9 atm and pNH3 = 10-9 atm, pH can be calculated as for a single gas. But this calculation is a little more complex because the solution is not particularly acidic and second dissociation constant of sulfurous acid becomes important. The value of pH obtained is about 5.8 showing that ammonia even at very low concentrations has an effect on pH.


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 – 7 x 105 cm-3; Hydroperoxy radical – 2 x 107 cm-3

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.


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.


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.


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 1011 cm-2 s-1 or about 0.7 x 1015 g (C) a-1. This is larger than the amount which arises from human activities (0.3 x 1015 g (C) a-1).


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 formation Free-energy 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.


N2O which is relatively stable in the troposphere, usually moves into stratosphere and undergoes following photochemical reactions:

N2O + hv ( 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 ( 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.


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.


In atmospheres of urban centres, under conditions of relatively low humidity, plenty of sunshine, a large amount of exhaust emissions from motor vehicles and moderate to low wind speeds, photochemical processes lead to a secondary pollution situation commonly termed “photochemical smog’. A large number of compounds and reactions have been characterized in the urban air where such smog situation occurs. The chemistry of this photochemical smog condition is extremely complex. Major photochemical processes associated with this condition have been discussed below.

1. Nitrogen oxide pseudo-equilibrium

The oxides of nitrogen, particularly NO and NO2 are at the root of photochemical smog problem. Oxidation of atmospheric nitrogen during high temperature combustion processes (particularly in motor vehicles) results in formation of NO which is further oxidized to NO2:

(a) O + N2 —-> NO + N

N + O2 —-> NO + O


N2 + O2 —-> 2NO

(b) 2NO + O2 ——> 2NO2

R1 = k1[NO]2[O2] where R1 and R2 are reaction rates and k1 and k2 are the rate

and, constants

NO + O3 —-> NO2 + O2

R2 = k2[NO][ O3]

The reaction of NO with oxygen at the concentrations found even in the polluted air is very slow, therefore, NO2 is mainly produced by oxidation of NO by ozone. In the polluted air, typical early morning concentrations of ozone and NO are 40 ppb and 80 ppb respectively. Values for k1 and k2 are 1.93 x 10-38 cm6s-1 and 1.8 x 10-14 cm3s-1 for ozone and NO respectively. From these values the calculation shows that R1 = 4.6 x 10-5 cm-3 s-1 and R2 = 3.8 x 1010 cm-3 s-1. This confirms the far greater importance of the oxidation of No by ozone.

The NO2 produced in this way can be photodissociated back to NO. Thus a sequence of reactions describing its destruction and regeneration can be given:

NO2 + hv ( O(3P) + NO

O(3P) + O2 + M ——> O3 + M

O3 + hv (300-330 nm) ——> O2 + O(1D)

O3 + HO2 ——> 2O2 + OH ( R = 1.1 x 10-14)

O3 + NO —–> NO2 + O2 (R = 2.3 x 10-12)

In a volume of air in steady-state where production and destruction rates of NO2 are equal and where oxidation of NO by oxygen is assumed to be unimportant, the reaction rate may be written as:

k2[NO][O3] = J[NO2]

where J is effective first-order rate constant for photodissociation. The equation may be rearranged as:

J/k2 = [NO][O3]/[NO2]

where the term on right-hand side may be ignored as a pseudo-equilibrium constant relating the partial pressures of NO, NO2 and O3. The value of J will varies with change in intensity of sunlight throughout the day. However, measurements have shown that overall the equality implied in this equation holds in the polluted atmosphere. During first half of the day radiation intensity increases which means J will increase and during this period increasing amounts of ozone and NO would be expected. Since both these are produced by destruction of NO2, the amount of ozone should approximately equal the amount of NO.

Measurements from polluted atmospheres show that neither of the above predictions are borne out. The level of NO rises in the early morning but level of ozone rises much more slowly. Further, the fact that the level of NO2 falls by mid-day is even more in contrast to the theoretical prediction. A possible explanation for these observations is that the observed rises and falls in the concentrations of pollutants are merely functions of the pattern of generation and dispersion in the atmosphere.

2. Role of organic molecules in smog

Under constant illumination the rise in the level of ozone indicates a decreasing NO:NO2 ratio in the pseudo-equilibrium. For the latter to happen, another source of oxidant is needed because above described sequence of reactions does not result in any overall production of ozone. However, ozone production in polluted atmosphere may be explained by following scheme.

As in unpolluted atmosphere, oxidation in polluted atmosphere also occurs through reactions in which hydroxyl radical plays a key role. The hydroxyl radical attacks a veriety of pollutants in the urban air resulting in formation of free radicals like methyl radical (CH3), acetyl (CH3CO) and atomic hydrogen (H) which may become involved in subsequent reactions which oxidize to NO to NO2 and regenerate hydroxyl radical at the same time.

(a) Alkanes in smog: Presence of alkanes such as methane in polluted air provides a way in which NO can be oxidized to NO2 without consuming ozone. For example, methane may be oxidized by OH radical to produce methyl radical which further undergoes a series of reactions:

CH4 + OH —-> H2O + CH3 (R = 2.4 x 10-12)

CH3 + O2 +M —-> CH3O2 + M

CH3O2 + NO —-> CH3O + NO2 (R = 7.0 x 10-12)

CH3O + O2 —–> HCHO + HO2 (R = 5.0 x 10-13)

HCHO = hv ( 2 H + CO

CO + OH —–> CO2 + H (R = 1.35 x 10-13)

HO2 + NO —–> NO2 + OH (R = 4.3 x 10-12)

HO2 radical can also react photochemically or with ozone, atomic hydrogen or atomic oxygen to regenerate OH radical. HCHO can photodissociate into atomic hydrogen or react with oxygen to give the HO2 radical and CO.

The above reactions can be summed up and show the importance of methane in generating NO2 in photochemical smog:

CH4 + 2 O2 + 2 NO —-> H2O + HCHO + 2 NO2

This indicates net oxidation of NO in a manner that has not used ozone, therefore, it is different from pseudo-equilibrium situation.

(b) Aldehydes in smog: Aldehydes also provide effective ways of oxidizing NO to NO2. For example, acetaldehyde is attacked by OH radical producing acetyl radical which undergoes following subsequent reactions:

CH3CO + O2 —–> CH3COO2

CH3COO2 + NO —–> NO2 + CH3CO2

CH3CO2 ——> CH3 + CO2

Methyl radical produced is oxidized as described above. There are analogous reactions for higher aldehydes.

(C) Atomic hydrogen in smog: The atomic hydrogen produced by attack of OH on CO or photodissociation of HCHO can react with HO2 radical to produce two OH radicals that can initiate further attack on organic compounds in air.

OH + CO —–> CO2 + H

H + HO2 —–> OH + OH

Atomic hydrogen can also form HO2 radical which can oxidize NO to NO2:

H + O2 + M —–> HO2 + M

HO2 + NO —–> NO2 + OH

In general, hydrocarbons present in the polluted urban air promote the oxidation of NO to NO2 by reactions of the types described above. The NO2 is subsequently photolysed to produce NO for reoxidation and increasing amount of ozone.

NO2 + hv ( O(3P) + NO

O(3P) + O2 + M —–> O3 + M

Though there are losses in the above described scheme, the built-up of ozone throughout the day can thus be well explained.

3. Other products in photochemical smog

A number of other features of photochemical smog can also be explained by photochemical mechanism described above.

1. Formation of PAN: Peroxyacetylnitrate (CH3COO2NO2) or PAN is a major eye-irritant found characteristically in photochemical smog. The peroxyacetyl radical (CH3COO2) produced by attack of acetyl radical on oxygen can combine with NO2 to form PAN:

CH3COO2 + NO2 ——> CH3COO2NO2

PAN is the principal member of a group of rather similar nitrated compounds which includes higher peroxyalkyl compounds such as peroxypropionyl nitrate which has also been detected in low concentrations in photochemical smog. There is also much current interest in the natural production of compounds like PAN.

2. Formation of N2O5 : NO2 is oxidized by ozone to NO3 which subsequently reacts with NO2 to form N2O5. The NO3 may also react with NO to produce more NO2.

NO2 + O3 —-> NO3 + O2

NO3 + NO2 ——> N2O5

NO3 + NO ——-> 2NO2

3. Formation of nitric acid: The OH radical formed in the smog reacts with NO to form HNO2 and with NO2 to produce HNO3. There may be reaction between NO and NO2 to form HNO2.

NO + OH —-> HNO2

NO2 + OH + M—–> HNO3 + M

NO + NO2 + H2O ——> 2HNO2

HNO2 undergoes photodissociation to produce NO and provide a source of OH radicals.

HNO2 + hv ( NO + OH

4. Formation of hydrogen peroxide:
Formaldehyde in polluted air is an important source of atomic hydrogen and hence OH and HO2 radicals:

HCHO + hv ( 2H + CO

H + O2 + M ——> HO2 + M

HO2 + NO —–> NO2 + OH

The OH and HO2 radicals may produce H2O2:

OH + OH + M —–> H2O2 + M

HO2 + HO2 ——–> H2O2 + O2 (R = 3.8 x 10-14)

5. Oxidation of sulfur dioxide: Sulfur dioxide can be oxidized under photochemical conditions but the S-O bond is very strong. So the sulfur dioxide can not undergo photodissociation as in the familiar case of NO2. The oxidation of SO2 involves OH radical:

OH + SO2 ——> HSO3

HSO3 + O2 —–> HSO5 or,

HSO5 ——-> HO2 + SO3 HSO3 + O2 ——> HO2 + SO3

SO3 + H2O —–> H2SO4

There is increasing evidence that the two middle reactions occur as a single reaction.

4. Degradation of larger organic molecules

Larger organic molecules (other than methane and acetaldehyde) are also split up in photochemical smog.

(a) Alkanes: Degradation of large alkane molecules (e.g. butane) starts with attack by OH radical:

OH + CH3CH2CH2CH3 ——-> H2O + CH3CH2CH2CH2

O2 + CH3CH2CH2CH2 ——-> CH3CH2CH2CH2O2


CH3CH2CH2CH2O + O2 ——-> CH3CH2CH2CHO + HO2

CH3CH2CH2CHO + hv ——–> CH3CH2CH2 + HCHO

(b) Alkenes: Large alkane molecules may be degraded by being attacked by ozone, atomic oxygen (O(3P) or OH radical. Attack by OH radical predominates in polluted atmosphere. A typical reaction scheme may be illustrated using butane as example:




The process goes on and on.

The above reaction schemes show degradation of larger organic molecules into smaller ones resulting in greater predominance of low molecular weight compounds in typical urban atmosphere with exhaust fumes of automobile.

5. Heterogeneous reactions in photochemical smog

Gas-phase photochemical reactions may lead to formation of aerosols in polluted urban atmosphere and these give rise to visual obscurity associated with smog condition. High opacity of smog gives an exaggerated impression of the amount of particulate material present yet it is estimated that as little as 5% of pollutants present in photochemical smog could be converted into suspended particulate materials. Various heterogeneous reactions could occur on the surface of these particles or in cloud or rain droplets associated with smog. The material forming condensed phase of smog may consist of both inorganic and organic substances.

(i) Inorganic substances: These include metal oxides and the salts of acids produced within urban air. The acids (particularly sulfuric and nitric acids) are usually present in association with solid particles or more probably as droplets due to their high affinity for water. The latter can react rapidly with atmospheric ammonia. The ammonium sulphate and ammonium nitrate produced are important aerosols that are main causes for the reduction of visibility that accompanies photochemical smog.

(ii) Organic solids: Relatively little is known of the reaction pathways that produce organic particulate materials in the polluted urban air. Nitrogen has been detected in rather unusual reduced oxidation states on particles in photochemical smog. This nitrogen is thought to be present as nitriles, amines or amides bound onto the surface of soot particles. By denoting the soot surface as S, the process may be written as:

S-OH + NH3 —-> S-ONH4

(a phenolic hydroxy ammonium complex)

S-ONH4 —–> S-ONH2 + H2O (at higher temperature)



S-COONH4 —–> S-COONH2 + H2O (at higher temperature)

S-COONH2 ——> S-CN + H2O

Most thoroughly studied heterogeneous reaction in the atmosphere Is the oxidation of sulfur dioxide in atmospheric liquid droplets by the ozone, hydrogen peroxide or oxygen in the presence of a transition metal ion catalyst. This oxidation reaction has been discussed earlier and may proceed much faster in polluted urban atmospheres than in unpolluted atmospheres because the concentrations of oxidants (H2O2 or O3) and metal ion catalysts may be much higher. Metal ions may, in particular, be leached from particulates that are added into the air through anthropogenic activities. Leaching of metals from ash may be particularly significant in their surface concentrations being enriched. High amounts of soluble metal ions have been observed in association with fly ashes from the combustion of refuge derived fuels. A further mechanism for increasing the rate of oxidation involves dissolution of materials such as calcium oxide which are present in high concentrations in coal fly ash making the droplet alkaline:

CaO + H2O —–> Ca2+ + 2OH-

This allows dissolution of larger amounts of sulfur dioxide and thus increases the rate of catalytic oxidation. Alternatively, dissolution of ammonia from a polluted atmosphere will also increase the pH and enhance both the dissolution and oxidation of sulfur dioxide.

Oxidation of sulfur dioxide may also occur via absorption of gas onto solid surfaces followed by subsequent oxidation. However, the surface area of particulate material even in polluted atmosphere is quite small and, therefore, such mechanism requires some method of ‘cleaning’ the surface in order to make oxidation process significant. If particulates are wet, this mechanism may be effective since water would ‘clean’ the surface of particulate material.

In the atmosphere, changes in the size and/or composition of particles also occur. These include leaching of particulate material by water, oxidation or reduction of particles. Zinc vapour from copper smelters condenses to form highly angular and crystalline zinc oxide crystals in the atmosphere. These are gradually degraded, then rounded and now acquire a carbonaceous coating. Slowly zinc oxide core decomposes and particle ends up as a carbonaceous pseudomorph with little or no zinc. Possibly, carbonaceous particles are formed by reduction of zinc oxide following deposition of hydrocarbons onto the surface of the particle.


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The atmosphere of Earth comprises air envelop over the Earth’s surface extending to several kilometers into the space. Air is a mixture of gases and is held to the Earth by gravita­tional attraction. The density of atmosphere is maximum at sea level and decreases rapidly upward.

The development of atmosphere is closely related to the geological and geochemical processes and to the activities of living organisms. Initially the Earth did not possess atmosphere and it was created later in the course of Earth’ evolution. Important components of atmosphere i.e. nitrogen, carbon dioxide and water vapor arose in course of volcanic activities which brought these from the depths of lithosphere. Atmospheric oxygen was added later in Earth’s history as a result of photosynthetic activity of living organisms. In its turn, atmosphere has exerted a great influence on the evolution of lithosphere. Throughout Earth’s history, atmosphere has played important role in physical and chemical erosion of rocks. Winds, atmospheric precipitation and changes in atmospher­ic temperature and other atmospheric factors have been major factors in physical erosion of rocks while atmospheric oxygen and moisture have been extremely important in chemical erosion of rocks.

Evolution of atmosphere has played significant role in the evolution of hydrosphere since water balance of water bodies is influenced directly by precipitation and evaporation processes governed by atmospheric characteristics. On the other hand atmos­pheric processes have also been influenced by the state of hydro­sphere, and especially by the state of the oceans. Generally, the evolution of atmosphere and hydrosphere represents a single process.

Atmospheric factors have also played important role in influencing the evolution of biosphere. In the beginning the atmosphere had no oxygen and only anaerobic organisms could evolve. These organisms evolved the mechanisms by which they could split up water and release oxygen. As the oxygen accumulated into the atmosphere, the earlier reducing atmosphere turned into an oxidizing one. Only after sufficient accumulation of oxygen in atmosphere, aerobic organisms could evolve. Thus evolu­tion of complex living organisms is linked to increase in oxygen content of atmosphere which lead to the development of aerobic processes needed for the energetics of higher organisms. Carbon dioxide content of atmosphere is most important factor in activi­ties of autotrophic plants. Changes in its content during the course of evolution of atmosphere have exerted great influence on the structure and Earth’s plant cover. Further, the nature and characteristics of biotic communities depends on climatic condi­tions which are mainly governed by atmospheric factors.

Thus the atmosphere is very important component of the global environment and the important atmospheric features with a view to the study of global environment are:

1. Atmospheric stratification

2. Atmospheric gases and aerosols

3. Air pressure, winds and global air circulation

4. Atmospheric moisture and precipitation


The Earth’s atmosphere shows quite well defined layers one above the other. These layers are defined mainly on the basis of temperature. Broadly, the pattern consists of three relatively warm layers near the surface between 50 and 60 km and above 120 km separated by two cold layers between 10 and 30 km and about 80 km. Major layers in the atmosphere are as discussed below.

1. Troposphere

It is the lowermost layer of atmosphere where weather phe­nomena and atmospheric turbulence are most marked. This layer contains about 75% of total molecular or gaseous mass of atmos­phere and virtually all the water vapor and aerosols. Temperature in this layer decreases with altitude at a mean rate of about 6.5o C per kilometer. At the top the troposphere is limited and separated from the next higher layer by a layer called tropo­pause.

Tropopause is a temperature inversion level i.e. it is a layer where a layer of relatively warm air is present above a layer of colder air. This inversion level acts as a lid over most of the top of troposphere. As a result troposphere is largely self-contained and convection in it is effectively limited. The altitude at which tropopause is present is not constant but seems to be correlated with sea-level temperature and pressure which are in turn related to factors of latitude, season and daily changes in surface temperature. The altitude of tropopause varies from about 16 km at equator where heating and vertical turbulence are greatest to only about 8 km at poles. Thus troposphere ex­tends from ground surface to the altitude of 8-10 km at high latitudes to 16-18 km in equatorial zone. Physical processes within the troposphere determine the changes in the weather and exert a major influence on the climat­ic conditions in different regions of our planet. These processes include the absorption of solar radiation; the formation of fluxes of long-wave radiation, which is dissipated into outer space (and which changes in the higher layers of the air); and the water exchange that is associated with the formation of clouds and with precipitation

2. Stratosphere

This layer extends from tropopause to the altitude of about 50 km and contains most of the atmospheric ozone. Peak density of ozone in stratosphere is approximately at 22 km altitude. The ozone in this layer is responsible for the absorption of ultra-violet wave-lengths of solar radiation. Due to very low density of air such absorption results in large increase in the tempera­ture within this layer. The stratospheric temperature fairly generally rises with height in summers while the thermal struc­ture of this layer in winters is more complex. Thus marked sea­sonal changes of temperature affect the stratosphere. At the top stratosphere is limited by stratopause which has highest tempera­ture which may exceed 0o C.

3. Mesosphere

This zone of atmosphere extends from stratopause upward to about 80 km altitude. The temperature in this layer again de­creases from about 0o C at stratopause level to average of -90o C at the top of mesosphere. Air pressure in this layer is very low, decreasing from about 1 mb at 50 km altitude to 0.01 mb at 90 km altitude. Temperature again begins to increase above 80 km alti­tude. The temperature inversion level at 80 km altitude i.e. at top of mesosphere is termed mesopause.

4. Thermosphere

Above the mesopause, the thermosphere is the zone of ex­tremely low atmospheric density. This layer extends from 80 km to 100 km altitude. Molecular and atomic nitrogen and oxygen are the main constituent of this zone. The temperature in thermosphere increases with altitude owing to absorption of extreme ultra-violet radiation of 0.125-0.205 um wave-length by molecular and atomic oxygen.

5. Ionosphere

This layer is the region of high electron density extending between altitudes of 100 km and 300 km. Above 100 km altitude, this atmospheric zone is increasingly affected by cosmic radia­tion, solar X-rays and ultra-violet radiation. These radiations cause ionization of oxygen atoms and nitrogen molecules separat­ing the electrons from them. The frequency of ionized particles continues to increase upward alongwith the increase of tempera­ture in ionosphere.

6. Exosphere and Magnetosphere

Above the ionosphere, ions of oxygen, hydrogen and helium form the tenuous atmosphere generally called exosphere. In this zone the natural gas laws cease to be valid. Since natural atoms of hydrogen and helium are low molecular weight atoms, the atoms of these two elements escape from this zone into outer space. The escaped hydrogen is continuously replenished by breakdown of water and methane molecules near the mesopause. The escaped helium is similarly replenished through its formation by the action of cosmic radiation on nitrogen and from slow breakdown of radioactive elements in Earth’s crust.

The magnetosphere is the zone containing only plasma of electrically conducting gases. Charged particles are concentrated in two bands at altitudes of about 3000 km and 16,000 km. These zones form the Van Allen radiation belts or plasmosphere. The behavior of plasma particles in magnetosphere and the ‘precipita­tion’ of high energy plasma particles into Earth’s atmosphere produces ionization in lower atmospheric layers.

The stratification of atmosphere into distinct layers having specific structural and functional properties is very important feature of atmosphere since it governs the heat budget of Earth-atmosphere system, air motion and meteorological conditions. All these together result in weather phenomena in the troposphere which in turn govern the spatial and temporal pattern of climates i.e. the distribution of different climatic regimes in different geographical regions and their seasonal variations on Earth.


1. Atmospheric gases

The air forming the atmosphere is a colourless, odorless mixture of gases. Normally air consists largely of nitrogen (78%) and oxygen (21%) by volume. Remaining 1 % of the volume of air is made of small quantities of several other gases among which ex­tremely important gas is carbon dioxide (0.03%) because it can absorb heat and thus has primary role in maintaining temperature of Earth-atmosphere system. In the zone of atmosphere extending from ground surface upto the altitude of about 50 km, all the gaseous constituents of air are perfectly diffused among one another so as to give the air definite physical qualities just as if it were a single gas.

The gaseous composition of atmosphere is characterized both by permanent and variable components. Apart from carbon dioxide, another extremely important variable component of atmosphere is water vapour. It is a colourless, odorless gaseous form of water which mixes perfectly with other gases of air. Most of the atmos­pheric water vapor is concentrated in the troposphere zone. Changes in water vapor content of atmosphere over space and time are determined by interaction between evaporation, condensation and horizontal movement of water in atmosphere. The degree to which water vapor is present in the atmosphere is designated as the humidity and is of tremendous importance in weather phenome­na. Condensation of atmospheric water vapor results in formation of clouds and fog while excessive condensation results in rain, storm, hail or sleet collectively termed precipitation. The atmospheric water vapor like carbon dioxide can absorb heat and, therefore, like carbon dioxide is extremely important in ‘green-house effect’ of atmosphere.

Ozone gas is also very important constituent of atmosphere because of its ability to absorb ultra-violet radiation. Most of the atmospheric ozone is concentrated in stratosphere. It is formed from oxygen in the atmosphere.

2. Atmospheric aerosols

Apart from gases atmosphere contains many types of extremely small and light matter suspended in it generally included under the term aerosol. The word aerosol includes a wide range of material that remains suspended for a period of time in the atmosphere and usually refers to small solid and liquid matter. Solid aerosols are usually defined as particles or particulates and are distinct from dust which includes large pieces of solid material (>0 m in diameter) which settle out of atmosphere due to gravitation after short period of suspension. While effects of dust are limited locally, smaller aerosols can be transported to long distances and affect air quality and climate on regional and global scales. Aerosols originate from two main sources and are accordingly termed primary aerosols or secondary aerosols.

(i) Primary aerosols: These include matter that has been swept into the atmosphere from the surface of Earth such as dry desert plains, lake beds and beaches, volcanic eruptions, forest fires, ocean surfaces, disintegration of meteors in atmosphere, biological sources (e.g. bacteria, pollen and (fungi) etc. About 90 percent of these aerosols are found in troposphere while they are also found in upper layers of atmosphere also. Primary aerosols of size 2.0-20.0 m are defined as coarse aerosols while those 2.0 m in diameter are defined as fine aerosols.

(ii) Secondary aerosols: These aerosols are formed after various types of chemical conversion processes in atmosphere which involve gases, other aerosols and atmospheric contents particularly the water vapor. Very little is know about the details of the chemistry of trace gases to aerosols. These aerosols are almost always less than 2.0 m in size at the time of their initial formation when they are at nucleation mode (<0.1 m) but grow rapidly to accumulation mode (upto 2.0 m). General age of a layer of these aerosols can be determined by the relative amount of nucleation versus accumulation sizes. The smaller aerosols coagulate rapidly and aerosols larger than accumulation mode are efficiently removed from atmosphere by wet and dry processes and depos­ited onto the Earth’s surface.

a) Sulfate aerosols: A large fractions of aerosols are sulfate aerosols. In the nucleation stage, liquid droplet of sulfuric acid grows rapidly to accumulation size and eventually forms a stable non-reactive particle containing sulfate. Most often eventual result is ammonium sulfate in ages aerosols or ammonium bisulphate. Typical concentrations of sulfate aerosols are :

Remote background area – 1-2 g/cubic meter

Non-urban continental areas – <10 g/cubic meter

Urban areas under anthropogenic influence – >10 g/cubic meter

b) Nitrate aerosols: Nitrate is another important component of aerosols and mainly comes from oxidation of nitrogen gas. Most common compound in fine aerosol range is ammonium nitrate. It is not as stable as ammonium sulfate and its concentration is con­trolled by the relative abundance of ammonium, nitrate, sulfate and the level of atmospheric temperature. Nitrate also exists in coarse aerosols as a reactive interchange between crustal ele­ments over the continents or sea salt (ammonium nitrate) over the ocean.

c) Other aerosols: Most other aerosols can be further classified into size components with their areas of impacts as given in the subsequnet Table-1.

Optical effects of aerosol particles

High concentration of particulate material in the atmosphere is responsible for the visible hazes. Suspended material can cause a range of rather unusual atmospheric phenomena such asblue moons, green suns and green flashes or arcs about the sun or moon.

The distances between aerosol particles are generally greater than 10-100 particle radii and with such distances, scattering of light by particles is incoherent. Therefore, optical effects due to atmospheric aerosol particles are explained by light scattering.

Table-1: Properties of miscellaneous aerosol particles present in atmosphere.


Class Size range m) Impact area


(i) Aerosol size

Aitken 0.005-0.1 Air electricity

Large 0.1-1.0 Suspended particulate

Giant 1.0-15.0 Suspended particulate

Dust >15.0 Gravitational fallout

(ii) Aerosol type

Small ions <0.001 Air electricity

Large ions 0.005-0.5 Atmospheric chemistry

Haze 0.08-2.0 Visibility, human respiratory problems

Mist & fog 1.0-20.0 Visibility, atmospheric chemistry

Cloud condensation

nuclei 0.05-5.0 Cloud processes

Main aerosol 0.5-5.0 Visibility,atmospheric mass chemistry, cloud processes, human respiratory problems


Reyleigh Law for unpolarized light applicable only to particles of radius <0.03 m implies that scattered intensity will be proportional to r6/4 where r = radius of particle and = wavelength of light. Blue colour of scattered light from sky is explained in terms of effective scattering at shorter wave-lengths as the scattered intensity is inverse function of wave-wavelength. Red colour of setting Sun is because light passes over a very long path through atmosphere and most of its blue region of spectrum is lost due to scattering. Spec­tacular sunsets after volcanic eruptions or bush-fires arise due to higher than normal concentrations of very fine particulate material in the atmosphere after such eruptions.


The atmosphere above the earth is not a static body of air. The air masses have different and definite patterns of movement and all such movements ultimately result in global air circula­tion which is most important phenomena from the point of view of climatic conditions in different regions, weather conditions and distribution of pollution. The global movement of air is inti­mately associated with changes in the air pressures at different places. Therefore, the present chapter deals with concepts asso­ciated with air pressure, winds and global air circulation.


The air being in gaseous state, is readily compressible and if we consider a vertical column of air, the air nearer the ground level is compressed more due to the greater weight of air mass above it. This greater compression of air nearer to ground surface results in greater air density. The higher density of air results in increased expandability i.e. greater air pressure. Thus we find a vertical distribution of air pressure; the air pressure is greatest at sea level and gradually decreases with height. This vertical distribution of air pressure is important in many aspects of atmospheric science. In addition of height, air pressure is also affected by temperature. Increase in temper­ature results in increase in air pressure. Since different re­gions of earth receive different amounts of solar radiation daily during different times and also yearly during different months, the air is heated to different levels in different regions. This results in different air pressures in different regions i.e. horizontal distribution of air pressure over the globe. This horizontal distribution of air pressure is particularly important from the point of view of the origin, direction and velocity of winds.

The air pressure is measured by barometer which may be mer­cury barometer or aneroid barometer. The unit representing air pressure is either inches or millimeter of mercury or milibar (mb) which is a unit of pressure equivalent to a force of 100 dynes per square centimeter. The standard air pressure at sea level is 29.92 inches = 760 mm of mercury or 1013 milibar (mb).

Isobaric maps

Pressure conditions can be shown on map by means of isobars, which are lines connecting all the places that have same baromet­ric pressure. On the daily weather map, which shows conditions for a specific time only, the isobars are essential in showing the location of moving centers of high or low air pressures. On climatic maps the isobars show average air pressures which have been computed from the data accumulated over the years. First attention shall be paid to the average world conditions of air pressures.

World pressure belts

Major air pressure belts found on earth’s globe are:

(a) Equatorial trough: In the general vicinity of equator, there is a broad zone of somewhat lower than normal pressure (1013 and 1009 mb) which is known as equatorial trough.

(b) Subtropical high-pressure belts: On the north and south of this equatorial trough, there are subtropical belts of high air pressures. These belts are centered on about latitudes 30 degrees North and South. In the Southern hemisphere, this belt is clearly defined. In the North­ern hemisphere, this belt is broken into two oceanic centers or cells, one over the eastern Pacific and other over the eastern North Atlantic. High pressure at these latitudes is the result of convergence of air at higher levels and is accompanied by a general subsidence of the air.

(c) Sub-polar low-pressure belts: Extending from latitudes 45 degrees North and South to ice-covered North and South polar centers respectively are two broad belts of low pressure. In Southern hemisphere, there is a well developed subpolar low-pressure belt extending over the continuous expanse of southern ocean. The low pressure in these high latitudes in both the hemispheres is the result of numerous intense storms, each of which is a moving low-air pressure cen­ter. The pressure belts shift seasonally through several degrees of latitude alongwith the isotherm belts accompanying them. These seasonal shifts are important in explaining the world climates.

Northern hemisphere pressure centers

In the Northern hemisphere, the belted arrangement typical of Southern hemisphere is absent. This is due to the powerful influence that the vast land masses of Northern America and Asia separated by North Atlantic and North Pacific oceans exert over the pressure conditions in the Northern hemisphere.

Land areas develop high-pressure centers at the same time when winter temperatures fall far below those of adjacent oceans.Land areas develop low-pressure centers in summers when land surface temperatures rise sharply above temperatures over the adjoining oceans. Ocean areas show centers of pressures opposite to those on the lands, as seen in the January and July isobaric maps. In winters, the pressure contrasts as well as the thermal contrasts are greater. Over north central Asia, there develops Siberian high with pressure average exceeding 1036 mb. Over the central North America, there develops a clearly defined but much less intense center of high pressure, called the Canadian high.

Over the oceans, there are found Aleutian low and Icelandic low, named after the localities over which they are centered. These two low-pressure areas have much cloudy, stormy weather in win­ter, whereas the continental highs characteristically have a large proportion of clear, dry days.

In summer, pressure conditions are exactly opposite of winter conditions. Asia and North America develop lows, but the low in Asia is more intense. It is centered in southern Asia where

it is fused with the equatorial low-pressure belt. There are two well developed cells of the subtropical belt of high pressure over the Atlantic and Pacific oceans. These high pressure cells are shifted northward of their winter position and are considera­bly expanded. These cells are termed Bermuda high and Hawaiian high respectively.


Wind is simply defined as air in motion. Local winds are produced on a local scale by processes of heating and cooling of lower air. Following two categories of local winds may be recog­nized.

(i) Katabatic winds: The first category includes local winds in hilly or moun­tainous regions, where on clear and clam nights, heat is rapidly lost by ground radiation. This produces a layer of cold, dense air close to ground. A component of the force of gravity, acting in the downslope direction, causes this cold air to move down the mountain sides, pouring like a liquid into ravines and thence down the grade of the larger valley floors. Mountain breezes of this origin are of a variety termed katabatic winds. Particular­ly strong, persistent katabatic winds are felt on the great ice caps of Greenland and Antarctica where the lower air layer becomes intensely chilled. Certain occurrences of severe blizzards in these regions are katabatic winds.

(ii) Convection winds: In the second category are included land and sea breezes, which affect only a coastal belt a few km in width. Heated during the day by ground radiation, the air over land becomes lighter and rises to higher elevations. Somewhat cooler air over the adjoining water then flows land-ward to replace the rising warmer air creating a pleasant sea breeze. At night, rapid cooling of the land results in cooler, denser air which descends and spreads seaward to create a land breeze. These daily alternations of air flow are parts of simple convection systems in which flow of air takes a circular pattern in vertical cross section. Land and sea breezes are limited to periods of generally warm, clear weather when regional wind flows is weak, but they form an important element of the summer climate along coasts.

Irrespective of whether there are pressure centers or belts, a pressure gradient always exists, running from higher to lower pressure. If isobars are closely placed, it indicates that the pressure gradient is strong and pressure changes occur rapidly within a short horizontal distance. Widely placed isobars indicate a weak pressure gradient. Most of the widespread and per­sistent winds of the earth are air movements set up in response to pressure differences. The pressure gradient force acts in the direction of pressure gradient and tends to start the air flow from higher to lower air pressure. Strong pressur gradients cause strong winds and vice versa. Calm exists in the centers of high pressures.

Coriolis force and geostrophic winds

If the earth did not rotate upon its axis, winds would follow the direction of pressure gradient. However, the rotation of earth upon its axis produces another force, the Coriolis force which tends to turn the flow of air. The direction of action of Coriolis force is stated in the Ferrels’s Law‚ which states that any object or fluid moving horizontally in the Northern hemi­sphere tends to be deflected to the right of its path of motion, regardless of the compass direction of the path. In the Southern hemisphere, similar deflection occurs towards the left of the path of motion. The Coriolis force is absent at the equator but increases progressively poleward. It should be noted especially that the compass direction is not of any consequence. If we face down the direction of motion, turning will always be towards the right hand in Northern hemisphere. Since the deflective force is very weak, it is normally apparent only in freely moving fluids such as air or water. Ocean currents patterns are, to some ex­tent, governed by it, and streams occasionally will show a tend­ency to undercut their right-hand banks in hemisphere. Driftwood floating in rivers at high latitudes in Northern hemi­sphere, concentrates along the right-hand edge of the stream.

Applying these principles to the relation of winds to pressure, the gradient force (acting in the direction of the pressure gradient) and the Coriolis force (acting to the right of the path of flow) reach a balance or equilibrium only when the wind has been turned to the point that it flows in the direction at right angles to the pressure gradient i.e. parallel with the isobars. The ideal wind in this state of balance with respect to the forces, is termed the geostrophic wind for cases in which the isobars are straight. In general, air flow at high altitudes parallels the isobars. The rule for the relation of wind to air pressure in the Northern hemisphere states that: Standing with back to the wind, the low pressure will be found on the left-hand side and high pressure on the right-hand side.

Between the ground level and altitude of about 2000-3000 ft., still another force modifies the direction of wind. This force is the friction of air with ground surface. This force acts in such a way as to counteract, in part, the Coriolis force and to prevent the wind from being deflected until parallel with isobars. Instead, the wind blows obliquely across the isobars, the angle being from 20 to 45 degrees.


The wind systems present on the earth’s surfaces may be categorized as following:

(1) Doldrums: In the equatorial trough of low pressure, intense solar heating causes the moist air to break into great convection columns, so that there is a general rise of air. This region, lying roughly between 5 degrees N and 5 degrees S latitudes was long known as the equatorial belt of variable winds and calms or the doldrums. There are no prevailing surface winds here, but a fair distribution of directions around the compass. Calms prevail as much as a third of the time. Violent thunderstorms with strong squall winds are common. Since this zone is located on a belt of low pressure, it has no strong pressure gradients to induce persistent flow of wind.

(2) Trade wind belts: In the north and south of the doldrums are the trade wind belts. These roughly cover the two zones lying between latitudes 5 degrees and 30 degrees N and S. These winds are the result of a pressure gradient from the subtropical belt of high pressure to the equatorial trough of low pressure. In the Northern hemi­sphere, air moving towards equator is deflected by the earth’s rotation to flow southwestward. Thus the prevailing wind is from the northeast and the winds are termed northeast trade winds. In the Southern hemisphere, deflection of moving air towards left causes the southeast trades. Trade winds have a high degree of steadiness and directional persistence. Most winds come from one quarter of the compass.

The systems of doldrums and trades shifts seasonally north and south, through several degrees of latitudes alongwith the pressure belts that cause them. Because of the large land areas of northern hemisphere, there is a tendency for these belts to be shifted farther north in summer (July) than they are shifted south in winter (January). The trades are best developed over Atlantic and Pacific oceans, but are upset in the Indian Ocean region due to proximity of the great Asian land mass.

(3) Winds of horse latitudes: Regions between latitudes 30 and 40 degrees in both hemi­spheres have long been called the subtropical belts of variable winds and clams or the horse latitudes. These coincide with subtropical high-pressure belts. However, these are not continu­ous belts and high-pressure areas are concentrated into distinct centers or cells located over the oceans. The apparent outward spiraling movement of air is directed equatorward into the east­erly trade wind system; poleward into the westerly trade wind system. The cells of high pressure are most strongly developed in the summer (January in Southern and July in Northern hemisphere). There is also a latitudinal shifting following the sun’s declina­tion. This amounts to less than 5 degrees in Southern hemisphere, but it is about 8 degrees for the strong Hawaiian high located in the north eastern Pacific.

Winds in these regions are distributed around a considerable range of compass directions. Calms prevail upto quarter of the time. The cells of high pressure have generally fair, clear weather, with a strong tendency to dryness. Most of the world’s great deserts lie in this zone and in the adjacent trade-wind belt. An explanation of the dry, clear weather lies in the fact that the high pressure cells are centers of descending air, settling from higher levels of the atmosphere and spreading out near the earth’s surface and the descending air becomes increas­ingly dry.

(4) Westerlies: Between the latitudes 35 and 60 degrees, both N and S, is the belt of westerlies or the prevailing westerly winds. Moving from the subtropical high-pressure centers towards the subpolar lows, these surface winds blow from a southwesterly quarter in the Northern hemisphere and from a northwesterly quarter in Southern hemisphere. This generalization is somewhat misleading because winds from polar direction are frequent and strong. More accurately, winds within the westerly wind belts blow from any direction of the compass but the westerly components are defi­nitely predominant. In these belts, storm winds are common cloudy days with continued precipitation are frequent. Weather is highly changeable.

In Northern hemisphere, land masses cause considerable disruption of the westerly wind belt but in Southern hemisphere, there is an almost unbroken belt of ocean between the latitudes 40 and 60 degrees S. Therefore, in Southern hemisphere the west­erlies gain great strength and persistence.

(5) Polar easterlies: The characteristic wind systems of the arctic and antarctic latitudes is described as polar easterlies. In the Antarctic, where an ice-capped mass rests squarely upon the south pole and is surrounded by a vast oceanic expanse, polar easterlies show an outward spiraling flow. Deflected to the left in Southern hemi­sphere, the radial winds would spiral counterclockwise, producing a system of southeasterly winds.


In the Northern hemisphere, continents of Asia and North America exert powerful control upon the conditions of atmospheric temperature and pressure. Since pressure conditions control winds, these areas obviously develop wind systems that are rela­tively independent of the belted system of earth’s surface winds which is very developed in the Southern hemisphere. These inde­pendent wind systems are termed monsoon winds.

In summer, southern Asia develops a center of low pressure, into which there is a considerable flow of air. This may be a heat low (thermal low)“ limited to the lower levels of atmosphere. Warm, humid air from the Indian ocean and southwestern Pacific moves northward and northwestward into Asia, passing over India, Indochina and China. This air flow is summer monsoon winds which is accompanied by heavy rainfall in southeast Asia.

In winter, Asia is dominated by a strong center of high pressure, from which there is an outward flow of air reversing that of the summer monsoon. This flow is the winter monsoon winds which blows southward and southeastward toward the equatorial oceans and brings dry, clear weather for a period of several months.

The North America is smaller in extent as compared to Asia, and so it does not have such remarkable extremes of monsoon winds as is experienced by southeast Asia. Nonetheless, North America also experiences an alternation of temperature and pressure conditions between winter and summer. In summer, there is a prevailing tendency for air originating in the Gulf of Mexico to move northward across central and eastern part of U.S.A. In winter, there is a prevailing tendency for air to move southward from sources in Canada.

The continent of Australia also shows a monsoon effect, but being situated south of equator, it exhibits conditions reverse to those in Asia.


The surface wind systems described above represent only a shallow basal air layer of a few thousand feet thickness, whereas the troposphere is five to twelve miles thick. Since 1945 much knowledge has been gained about the nature of air flow at higher levels in troposphere and weather maps of upper air conditions have been drawn. It has been found that high above, there are slowly moving high- and low-pressure systems but these are gener­ally simple in pattern with smoothly curved isobars. Winds, which may be extremely strong and follow the isobars closely, move counterclockwise around the lows (Northern hemisphere), but clockwise around the highs. In general or average pattern of upper air flow, two systems dominate:

(i) Westerlies: This is the system of winds blowing in a com­plete circuit above the earth from about latitude 20 degree almost to the poles in both hemispheres. At high latitudes these westerlies constitute a circumpolar whirl, coinciding with a great polar low pressure center. Towards low latitudes the pressure rises steadily at a given altitude, to form two high-pres­sure ridges at latitudes 15 to 20 degrees N and S. These are the high altitude parts of the subtropical highs, but are shifted somewhat equatorward. In the high-pressure zones, wind velocities are low, just as in the horse latitudes at sea level.

(ii) Equatorial easterlies: This second major global air circu­lation system is between the high-pressure ridges where there is a trough of weak low-pressure, in which the winds are easterly. At lower elevation their influence spreads into somewhat higher latitudes as the trade winds.


The upper-air westerlies tend to form somewhat serpentine and meandering paths, giving rise to slowly moving upper air waves‚ in which the winds are turned first equatorward, and then poleward. At altitudes of 30,000 to 40,000 feet, associated with the development of such upper air waves, are narrow zones in which wind streams attain velocities upto 200 to 250 miles per hour. This phenomenon is termed the jet stream and it consist of pulse-like movements of air following a broadly curving track. In cross section, the jet may be likened to a stream of water moving through a hose, the center line of highest velocity being surrounded by concentric zones of less rapidly moving fluid.

Most important function of upper air waves is that by means of these, the warm air of tropics is carried far north at the same time that cold air of polar regions is brought equatorward. In this way the horizontal mixing i.e. advection develops on a gigantic scale and serves to provide heat exchange between re­gions of high and low insolation.


The water evaporating from the surface of water bodies and also transpiring from the plants is held in the atmosphere as atmospheric moisture. The amount of water vapor held in the atmosphere at a given time varies widely from place to place. It ranges from virtually nothing in cold, dry air of arctic regions in winter to as much as 4 to 5 percent of the volume of atmos­phere in humid, hot tropical areas. The atmospheric water vapor returns to Earth’s surface in the form of precipitation which includes rain, snow, sleet and hail. This cycle of water from Earth’s surface to atmosphere and back to Earth constitutes very important part of global hydrological cycle. Further, this part of hydrological cycle is most important in creating and maintain­ing particular climatic conditions in different areas of Earth.

Important concepts related to atmospheric moisture and precipita­tion are discussed below.

Atmospheric humidity

The term humidity refers to the quantity of water vapor present in the air. For a given temperature, there is a definite limit to the quantity of moisture that can be held by the air. This limit is called saturation point. The actual total amount of water vapor present at a given temperature in a given volume of air is termed absolute humidity. The quantity of water vapor that can be held in a given volume of air increases with temperature i.e. absolute humidity is directly proportional to temperature.

The proportion of water vapor present relative to the maxi­mum quantity of water vapor that can be held at a given tempera­ture expressed as percentage is termed relative humidity at that temperature. At saturation point, relative humidity is 100%. Change in relative humidity can be caused in two ways:

(i) Addition of water vapor: When evaporation or transpiration adds water vapor to atmosphere, relative humidity increases. However, it is a slow process requiring that the water vapour diffuse upward through the air.

(ii) Decrease in temperature: Relative humidity can increase even without addition of water vapor to the air through a decrease in temperature because capacity of air to hold water vapor increases with decrease in temperature.

Dew point is that critical temperature at which air is fully saturated with the amount of water vapor present in it. If tem­perature falls below this point, condensation of atmospheric water vapor normally occurs.

When air rises or sinks in elevation, it undergoes changes of volume i.e. volume of air increases when it rise to higher elevation due to fall in atmospheric pressure and decreases when air sinks to lower elevation due to rise in atmospheric pressure. Due to this change in air volume with change in elevation, the value of absolute humidity can not remain a constant figure for the same body of air. Therefore, meteorology makes use of specif­ic humidity which is the ratio of weight of water vapor to weight of moist air (including water vapor). When a given air parcel rises or sinks in elevation without gain or loss of water vapor, the specific humidity remains constant. Specific humidity is often used to describe the moisture characteristics of a large mass of air. Its value ranges from 0.2 gm/kg for extremely cold, dry air over arctic regions in winter to 18-20 gm/kg for extreme­ly warm, moist air of tropical regions.

Another index of moisture used in meteorology is mixing ratio which is the weight of water vapor to weight of water vapor of dry air (excluding water vapor) stated in units of grams per kilogram. Mixing ratio commonly differs very little in actual numerical value of specific humidity.


When large masses of air are experiencing steady drop in temperature below dew point, condensation of water vapor occurs very rapidly within the clouds in atmosphere and precipitation occurs. This condition can not be brought about by simple cooling of air through loss of heat by radiation during night. Instead, rise of air to higher elevation is necessary. When air rises to higher elevation, its temperature drops even without loss of heat energy to outside. Because of drop of atmospheric pressure and increase in volume, air molecules strike each other less fre­quently and this imparts lower sensible temperature to air. In absence of condensation, the rate of drop of temperature is termed dry adiabatic rate and is about 5.5 degrees F per 1000 feet of vertical elevation. If water vapor in the air is condens­ing, latent heat is liberated which counteracts the temperature loss. The modified rate of adiabatic temperature loss in such condition, termed wet or saturation adiabatic rate becomes slightly lower to about 3 degree F per 1000 feet.

Water vapor does not necessarily condense in clean air even when the vapor pressure of water is many times greater than that required to form liquid water. The reason of this supersaturation condition of clean air is that the equilibrium vapor pressure over small droplets is much greater than that over plane surfaces (p ). Relation between radius of water vapor droplet and partial pressure of water vapor (pr ) above it is given by:

loge (pr /p ) = 2 M / L RTr

where, = surface tension of water (N/m); L = density of water (kg/cubic meter); M= molecular weight of water; R = universal gas constant (J/K/mol); T= absolute temperature in K; r = radius of droplet (m).

Presence of salts in water has pronounced effect on equilib­rium of relative humidity with a small water droplet .

Very small droplet of pure water requires that air be supersatu­rated for condensation to occur. Presence of even a small amount of salt lowers the water vapor pressure considerably and thus salt acts as condensation nucleus lowering the equilibrium rela­tive humidity remarkably. Removal of water vapor by condensation leads to fall in vapor pressure to below saturation level. In the cloud, the condition of supersaturation is maintained due to cooling of air rising to higher elevation. With cooling the relative humidity of air increases i,e, colder air becomes saturated at a lower water vapor content than warmer air, so the condition of supersatura­tion is maintained. Very high in a cloud, temperature may drop to below freezing point of water and so ice, snow, rain and hail can form.

Prior to actual precipitation clouds are formed in the atmosphere. These consist of tiny droplets of water 20-60 microns in diameter or minute crystals of ice. These are sustained in the atmosphere by the slightest upward movement of air. For the formation of cloud droplets, it is necessary that microscopic dust particles serve as condensation nuclei. Hydroscopic particu­late aerosols found in atmosphere serve as condensation nuclei. Precipitation occurs when condensation is occurring very rapidly in the clouds.


Formation of clouds and precipitation occurs only when large air masses rise to higher elevation. This rise of air masses occurs in three ways and accordingly precipitation is of three types:

(i) Convectional precipitation: Such precipitation results from a convectional cell which is an updraft of warmer air rising up because it is lighter than surrounding colder air. When bare land surfaces rapidly heated, it transmits radiant heat to overlying air. Air over warmer land is heated more than adjacent air and begins to rise in a tall column called thermal. As rising air cools adiabatically, it eventually reaches same temperature as the surrounding air and comes to rest. However, before coming to rest it may be cooled below the dew point and immediately conden­sation begins. The rising air column appears as cumulus cloud whose flat base shows the critical level above which condensation is occurring. Bulging ‘cauliflower’ top of this cloud represents the top of rising warm air column pushing into higher levels of atmosphere. If this convection column continues to develop, the cloud may grow into a cumulonimbus cloud mass from which heavy rainfall will occur. In most of the natural conditions, the unequal heating of ground serves only as a trigger effect to release a spontaneous updraft of air mass. Later it rises due to heating by release of latent heat from condensing water vapor. This heating causes air to continuously rise upward even during condensation. Such air is described as unstable air.

(ii) Orographic precipitation: This type of precipitation is related to mountains. Prevailing winds or other moving air masses may be forced to move over mountain ranges in some areas. As the air rises on the windward side of the range, it is cooled adia­batically. If cooling is sufficient, precipitation results on the windward side of range. Much orographic rainfall is actually of convectional type, in that it takes the form of heavy convection­al showers and storms. Storms are induced, however, by the forced ascent of unstable air as it passes over the mountain barrier.

(iii) Cyclonic precipitation: This type of precipitation occurs when air converges in cyclonic storms or eastward-moving centers of low pressure and is forced to rise resulting in cooling and condensation. Much of the precipitation in middle and high lati­tudes is of such type.