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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