Tuesday, March 11, 2008
Chemistry of Atmosphere
GAS-PHASE ATMOSPHERIC CHEMISTRY
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}
Since,
d[A]/[A]t = -kdt
or,
d[A]/dt = -k[A]t
so,
Fo = [A]/ = k[A]
or,
= 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.
GAS SOLUBILITY IN ATMOSPHERIC WATER
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-1 atm-1)
Hydrogen peroxide 2 x 105
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 hydrolyses 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:
T(HCHO) = KH pHCHO(1+K)
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 2.7 x 10-2 10-7
Ammonia 90 1.6 x 10-5 -
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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.
DISSOLUTION OF GASES IN SUSPENDED DROPLETS
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(-Dn22t2/r2)
n=1
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 50m 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.
CHEMISTRY INSIDE ATMOSPHERIC DROPLET
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)
or,
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.
CHEMICAL PROCESSES ON AEROSOLS
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,
hv
2CO + O2 ----------------> 2CO2
TiO2, ZnO
hv
2N2 + 6H2O -------------> 4NH3 + 3O2
TiO2
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.
RAINWATER CHEMISTRY
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.
ATMOSPHERIC PHOTOCHEMISTRY
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.
1. REACTION WITH SULPHIDES
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.
2. REACTIONS WITH AMMONIA
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
3. REACTION WITH CARBON MONOXIDE
Hydroxyl radical on reaction with carbon monoxide yields carbon dioxide and hydrogen radical.
CO + OH ----------> CO2 + H
4. REACTION WITH FORMALDEHYDE
Formaldehyde found in trace quantities and formed in various atmospheric reactions is oxidized by OH radical in the following manner:
HCHO + OH --------> HCO + H2O
5. REACTION WITH METHANE
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
6. REACTION AMONG HYDROXYL HYDROPEROXY RADICALS
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.
HYDROXYL AND HYDROPEROXY RADICALS AND PHOTOCHEMICAL REACTIONS IN ATMOSPHERIC DROPLETS
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.
1. OXIDATION REACTIONS INVOLVING HYDROXY RADICAL
(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
2. REACTIONS INVOLVING HYDROPEROXY RADICAL
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)
3. OXIDATION OF NATURALLY OCCURRING ALDEHYDES AND ALCOHOLS
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.
CHEMISTRY OF IONOSPHERE
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 CYCLE IN ATMOSPHERE
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).
OZONE CHEMISTRY OF ATMOSPHERE
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
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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 (<175 nm) ------> O(3P) + O(1D)
Oxygen atom in excited state on collision with some diatomic molecule (M2) yields oxygen atom in ground state i.e. O(3P):
Collisional
O(1D) + M2 -----------------> O(3P) + M2
deactivation
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.
NITROGEN CHEMISTRY OF STRATOSPHERE
N2O which is relatively stable in the troposphere, usually moves into stratosphere and undergoes following photochemical reactions:
N2O + hv (<260 nm) --------------> N2 + O(1D)
N2O + O(1D) ----------> N2 + O2
N2O + O(1D) ----------> 2NO
NO or NO2 which may also move from troposhpere into stratosphere or are produced in the stratosphere, undergo following reactions there:
NO + O + M ------> NO2 + M
NO2 + OH + M -----> HNO3 + M
HNO3 + hv (<330 nm) ------> OH + NO2
Ozone layer in stratosphere absorbs sufficient amount of UV radiation so that at tropopause HNO3 has photochemical lifetime of about 10 days. This time is long enough for much of it to cross the tropopause and come down with rainfall thus being removed from stratosphere.
CHLORINE CHEMISTRY OF 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.
CHEMISTRY OF PHOTOCHEMICAL SMOG
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 (<310 nm) ----> 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 (<360 nm) -----> 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 (<310 nm) ----> 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 (<400 nm) -------> 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 (<370 nm) ---------> 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
CH3CH2CH2CH2O2 + NO ------> CH3CH2CH2CH2O + NO2
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:
OH + CH3CH=CHCH3 --------> CH3CHOHCHCH3
O2 + CH3CHOHCHCH3 -------> CH3CHOHCH(O2)CH3
NO + CH3CHOHCH(O2)CH3 -------> CH3CHOH + CH3CHOH + NO2
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)
(amine)
S-COOH + NH3 -----> S-COONH4
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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Cycling of matter in environment
CARBON, NITROGEN AND SULFUR IN ENVIRONMENT
In the global environment, a vast number of elements exist in a variety of chemical species and are continually transformed from one species to another. These transformations from one chemical species to another involve cycling of these elements or chemical species amongst different components of the environment i.e. amongst atmosphere, lithosphere, hydrosphere and biosphere. These cycles of elements involving different components of the environment are, therefore, considered as biogeochemical cycles. These biogeochemical cycles are highly complex and interact strongly with each other and, therefore, are of fundamental importance in maintaining global environmental balance and in understanding the dynamics of environment. Further, human activities cause increase or decrease in natural amounts of chemical species in the environment or cause addition of chemical species not found in nature resulting in disturbance in natural biogeochemical cycles. Such disturbances constitute environmental pollution, which has profound impact on the stability of global environment. The biogeochemical cycles of element carbon, nitrogen and sulfur are most important from the point of view of global environmental balance. Therefore, important features of the cycles of these three elements have been discussed below.
CARBON CYCLE
The carbon cycle is mainly associated with living matter, although inorganic carbon provides important segments to complete the cycle. The cycling of carbon is strongly controlled by its storage in natural reservoirs. The time period of such storage may range from millennia in rocks, through decades in deep ocean layers to seasons in active biota. Relevant time periods of such storage suggested by Warneck (1988) are:
1. Geological activity involving rocks: 2,400 to 30,000 years
2. Soil humus: 200 years
3. Long-term biosphere storage: 75 years
4. Short-term biosphere storage: 15 years
5. Ocean mixed layers: 4 to 10 years
Estimates of mass content of carbon in various global reservoirs are given in the Table-1.
Carbon in oceans
Major storage of carbon in oceans occurs in the intermediate and deep water below the thermocline. The deep layers of oceans have a very slow mixing period and carbon remains in situ for atleast 20 years in these layers. Far above in oceans, in the mixed layer, which provides the main medium of interchange with the atmosphere, carbon storage is about 1.5 orders of magnitude lower. Ninety percent of the carbon in the oceans is stored as bicarbonate (CO32-) and about 9% as carbonate (CO3-). About 3% of carbon is present in organic matter in environment.
The mixing layer in oceans, broadly the layer above the thermocline, is assumed to be at depth of 75 meters. The average concentration of carbon dioxide in the oceanic surface layer (above the mixing layer) is 2.05 mmol m-3. This concentration rises rapidly with depth to about 2.29 mmol m-3 at the depth of about one-kilometer and remains fairly constant thereafter. The average oceanic carbon dioxide concentration is calculated to be about 2.25 mmol m-3. Since colder ocean water is able to hold more carbon dioxide, variations in its concentration occur with temperature of ocean water. The mass of carbon dioxide in the mixed layer is about the same as that in the atmosphere, with a total exchange between the two estimated to occur over a period of about seven years.
Table-1. Mass content of carbon in global reservoirs
| Reservoir | Carbon-content in Pg (1015 g) |
| OCEANS |
|
| 1. Total dissolved CO2 | 37400.0 |
| 2. Dissolved CO2 in mixed layer (75 m depth) | 670.0 |
| 3. Living biomass carbon | 3.0 |
| 4. Dissolved organic carbon | 1000.0 |
| SEDIMENTS |
|
| 1. Continental and shelf carbonates | 270 x 105 |
| 2. Carbonates in oceans | 230 x 105 |
| 3. Continental & shelf organic carbon | 100 x 105 |
| 4. Organic carbon in oceans | 200 x 104 |
| BIOSPHERE |
|
| 1. Terrestrial biomass | 650.0 |
| 2. Soil organic | 2000.0 |
| 3. Oceanic organic | 1000.0 |
| ATMOSPHERE (mostly as CO2) |
|
| 1. Pre-industrial estimate (290 ppmv) | 615.0 |
| 2. present estimate (350 ppmv) | 734.0 |
Organic carbon in oceans comes from precipitated remains of living organisms. About 80% of the precipitated material may be redissolved in the deep ocean layers. Dissolved organic carbon content of ocean waters is roughly estimated to be about 0.7 g m-3. Rest of the carbon in the ocean is particulate, mainly as calcium carbonate and this portion of oceanic carbon has a concentration of about 20 mg m-3. Living organisms contribute a total of only 3 Pg to the oceanic carbon storage.
Carbon in sediments and rocks
Carbon makes up only 0.032% of the Earth’s crust by mass. In terrestrial rocks, it is dissolved by rains or surface water over long periods of time and is carried by the surface runoff water to be deposited on the continental shelf sediments. In deeper oceans, deposits from organisms are built up on the ocean floor over millennia. Exchange of carbon from these locations occurs over thousands of years and is associated with activity of Earth’s crust. About two third of this carbon is inorganic carbon and rest is organic in form. Only about 1% of carbon in the form of oil and coal present in Earth’s crust can be used economically. It is estimated that if all the carbon stored in sediments is released suddenly, the atmospheric pressure will rise by 38 bars and the Earth’s atmosphere will become similar to that of planet Venus.
Carbon in biosphere
In the biosphere carbon is exchanged through:
Photosynthetic activity of photosynthetic living organisms, mainly the green plants
Release of carbon on decay and decomposition of dead living organisms
Respiratory activity of all the aerobic living organisms including both plants and animals
Release of carbon from soil humus
The mass of carbon is about three times higher than in living biosphere. The biospheric exchange processes are relatively inactive and the carbon storage may occur for 200 years. Long-lived species, particularly the plants store about 75% of the carbon present in the living biota. The major impacts on global carbon content present in the active biosphere occur in the forests, which store over 80% of the world’s biomass. Though estimates are uncertain because global distribution of different ecosystems is not known accurately, it is quite clear that tropical rain-forests, boreal forests and temperate forests are the most important ecosystems regarding storage and exchange of carbon.
Carbon in atmosphere
Exchange of carbon with the atmosphere occurs mainly through the biosphere with oceanic mixed layer being an important secondary source. Most important atmospheric form of carbon is CO2 gas and global estimates of its exchange between atmosphere and biosphere are:
1. Assimilation of CO2 into plants: 113 Pg Y-1
2. Re-release into atmosphere from:
Respiration of living organisms: 55 Pg Y-1
Microbial decay: 42 Pg Y-1
Soil humus: 10 Pg Y-1
Forest fires and agricultural burning: 1 Pg Y-1
3. Herbivore consumption: 5 Pg Y-1
In addition to CO2, other minor gases in the carbon chain are carbon monoxide (CO), methane (CH4) and non-methane hydrocarbons (NMHCs e.g. HCHO). Carbon dioxide gas is relatively inert while others are quite active in global atmospheric chemistry. Important features of atmospheric carbon species are discussed below.
1. Carbon dioxide: Though CO2 is a minor gas in the atmosphere in comparison with oxygen and nitrogen, it has major impact on global heat balance because of its high capacity of absorbing infra-red radiation. Continuously rising concentration of atmospheric CO2 due to various human activities, particularly the fossil-fuel burning, is major factor in global greenhouse warming. Anthropogenic carbon contributes about 3% of annual carbon loading. Further, its importance in relation to biosphere is supreme since it is required for photosynthesis and existence of biosphere depends on photosynthesis.
2. Carbon monoxide: About 90% of CO originates during photochemical production of methane in atmosphere. Some CO is produced during biomass burning and some during atmospheric oxidation of organic gases that are emitted from vegetation. Highest concentrations of CO are found in middle and high latitudes of Northern Hemisphere, which may reach 150 - 200 ppbv. The concentrations of atmospheric CO show a definite seasonal rhythm and are higher in summers than in winters. In Southern Hemisphere, CO concentrations are lower than in Northern Hemisphere by a factor of upto three. CO is removed from the atmosphere mainly by being oxidized to CO2.
3. Methane: This is a trace gas in atmosphere and is released mainly from rice paddies, wetland areas, enteric fermentation from animals and biomass burning. It has a uniform latitudinal distribution with an average concentration of about 1.6 ppmv. Major sinks of methane are temperate and tropical soils and oxidation to carbon monoxide.
4. NMHCs: This group includes a complex set of hydrocarbons with highly varying characteristics. Most of these are chemically active and have short lifetimes. The usual concentrations in the atmosphere are only few ppbv with localized peaks occurring near the sources. These compounds are removed from atmosphere usually by atmospheric photochemical reactions.
5. Particulate organic carbon (POCs): These complex mixtures of hydrocarbons, alcohols, esters and organics in particulate form. These are usually produced from secondary reactions (gas to particle conversions) and are important in cloud and precipitation processes. The concentrations of POCs in marine air may be around 0.1 to 0.5 g m-3 and in background continental air may be around 1.0 g m-3. In general, the composition of POCs has about 60% neutral compounds, 30% acids and 10% bases.
6. Elemental carbon: This comes into the atmosphere exclusively form biomass and fossil-fuel combustion. Its typical atmospheric concentration over continents is 0.02 g m-3. It is present as fine black powder and can be used as excellent tracer substance for studying long-range transport phenomena in atmosphere.
In addition to above forms, carbon is also present in the atmosphere as carbonyl sulfide, carbon disulfide and dimethyl sulfide. These compounds are important in sulfur-loading of atmosphere and have been discussed with atmospheric sulfur.
Table-2: Indicative characteristics of primary carbon compounds in atmosphere.
| Compoud | Major sources | Production (Tg Y-1) | Background concentration | Polluted concentration | Lifetime | Sinks |
| CO2 | Oceans, biosphere, fossil fuels | 7.6 x 104 | 350 ppmv | 380 ppmv | 5 years | Oceans |
| CO | Biomass burning, atmospheric photochemistry | 660.0 | <50 ppbv | 150-200 ppbv | 1-2 months | Oxidation to CO2 |
| CH4 | Animals, wetlands, decay of vegetation | 610.0 | 1650 pptv | >1800 pptv | 10 years | Oxidation to CO, soils |
| NMHCs | Vegetation, human activities | Variable | few ppbv | Variable | Variable | Photochemical reactions |
| POCs | Secondary atmospheric photochemistry | Small | 0.1 g m-3 | >2.0 g m-3 | 1 week | Wet and dry deposition |
| Elemental carbon | Biomass burning | Small | 0.2 g m-3 | >1.0 g m-3 | 1 week | Wet and dry deposition
|
NITROGEN CYCLE
Nitrogen is primarily exchanged between atmosphere, biosphere and soil. Following Table-3 shows the estimated total stored in the atmosphere and surface locations on a global scale.
Nitrogen in hydrosphere
In comparison to biosphere or atmosphere, very little nitrogen is present in oceans and continental surface waters. Over 95% of nitrogen stored in oceans is present in inactive molecular form. Only nitrate (about 2.5% of total oceanic nitrogen) and organic matter (about 1.5% of total oceanic nitrogen) have some active role. Oceanic nitrogen comes through river runoff from continents and wet and dry deposition from atmosphere. Its loss occurs through deposition to sediments in the bottom of oceans and through release to atmosphere in areas of biological activity. Nitrogen content in ocean water can vary spatially; for example, ammonia in surface oceanic waters varies between 0.05 to 2.0 mmol m-3 with smallest concentrations in the open oceans where biological activity is lowest. The amount of nitrogen released from oceans to the atmosphere (about 0.5 Tg Y-1) is quite low in comparison to that from other sources.
Table-3. Nitrogen storage in various
components of global environment.
| Location | Nitrogen storage in Tg (1012 g) |
| Lithosphere | 2 to 6 x 106 |
| Soil | 85 x 103 |
| Continental biomass | 10 x 103 |
| Atmosphere | 3.8 x 103 |
| Surface litter | 1.5 x 103 |
| marine biomass | 380.0 |
| Oceans | 23.0 |
| Human beings | 5.5 |
Nitrogen in rocks
The amount of nitrogen stored in lithosphere is much greater than the amounts stored in all other locations combined together. In lithosphere, most of the nitrogen is stored in primary igneous rocks and thus is not available to ecosystem. Weathering and other natural processes release only a very small fraction (<<1%) of this stored nitrogen into global ecosystem.
Nitrogen in soil and biosphere
Major active zone of nitrogen use and transfer occurs in the soil and biosphere on continents with very minor activity occurring in aquatic ecosystems. Inactive N2 of atmosphere is converted to form available to ecosystem through the process of nitrogen fixation, which mainly involves bacterial activities (though some nitrogen fixation also occurs during atmospheric lightening). Fixed nitrogen is made available first to plants in the ecosystem through mineralization to ammonia or through oxidation of reduced ammonia to nitrate (NO3-). This process termed nitrification occurs under aerobic conditions. The oxidized nitrogen in soil is returned to atmosphere through the process termed denitrification under anaerobic conditions.
Nitrogen content of soil determines the nitrogen availability to biosphere and various soil types differ in their nitrogen content. Most of the soils contain about 0.05% to 0.2% nitrogen by weight though richest organic soils may contain upto 0.5% of total mass as nitrogen. During rains, some of the soil nitrogen is leached by runoff or infiltration and reaches groundwater or river water to be transported elsewhere.
Nitrogen entering the plants mainly as nitrate or ammonium is assimilated there into a variety of organic nitrogenous compounds, mainly the proteins and amino acids which are passed on from plants to animals as food. Nitrogen then traverses to different trophic levels in the ecosystem as different animals eat each other. Finally, nitrogen is returned back to soil or atmosphere from the biosphere after death and decay of plants and animals. In the ecosystem, aerobic processes form NO2 also while anaerobic processes produce NO, N2O and N2. Most of these products is released to atmosphere.
All the processes and pathways involved with nitrogen cycle depend on the environmental conditions such as soil pH, water content, soil type etc. Temperature is crucial factor in nitrogen cycle because biological activity is highly sensitive to temperature.
Though nitrogen fixation is the natural source of biospheric nitrogen, nitrogen fertilizers added to soil and surface deposition of nitrogenous materials that are emitted into atmosphere by human activities have also become important inputs to biospheric nitrogen.
Nitrogen in atmosphere
Nitrogenous species important in global nitrogen cycle found in atmosphere are:
1. Molecular nitrogen: The N2 gas constitutes about 79% of air by volume and it provides the main source of nitrogen to biosphere through nitrogen fixation as discussed above.
2. Ammonia and ammonium: Ammonia is very important component of nitrogen cycle as it is the only water-soluble gaseous nitrogen species. It can directly act as plant nutrient being converted to ammonium (NH4+) which forms the atmospheric nitrogen aerosol component. About 54 Tg nitrogen is emitted to atmosphere per year and ammonia released from animal urea makes up about half of this. Nitrogen inputs through biomass burning depend on the nitrogen content of the biomass which differs in different ecosystems. Average nitrogen content of tropical forest wood is 0.45%, of tropical litter is 0.85%, of coniferous and deciduous forest wood is 0.32%, of fuel wood is 0.2% and of tropical grasses is 0.2% to 0.6%. Other minor sources include coal combustion, human excreta and fertilizers.
It is difficult to establish the global representative concentrations of ammonia and ammonium. Ammonia concentration is lowest over remote oceans (about 0.1 ppbv); while in continental background air it is 6-10 ppbv. The ammonia concentrations are higher in summers than in winters and during daytime than in night due to higher temperatures influencing the activities of soil-based microbial sources. The lifetime of ammonia is only about 6 days and so it is rapidly converted to ammonium, which is the major component of two most prevalent atmospheric aerosols, ammonium sulfate and ammonium nitrate. Concentrations of both these aerosols and the gas decrease exponentially with altitude. Major sink of these aerosols is wet and dry deposition that removes about 49 Tg of nitrogen per year from atmosphere.
3. Nitrous oxides: Apart from N2, nitrous oxide (N2O) is the other inert gas in the atmosphere. Its lifetime is about 179 years and its major sink is photochemical reactions in stratosphere. It is also a greenhouse gas. Major sources of N2O emission are soil and oceans through microbial processes. Highest concentrations of the gas over oceans occur in areas where strong upwelling brings deep-water nutrients to the surface waters. Emissions due to human activities are adding about 8% of the natural input. N2O emissions increase with higher temperature and moisture and, therefore, reach a daily maximum around noon and seasonal maximum in summers. Emissions can be greatly increased on a local scale by irrigation practices. The gas shows very little variation in global distribution due to its long lifetime and major natural sources. Depending on the photochemical activity, the concentration of gas decreases slightly with altitude in the troposphere.
4. Nitrogen oxide species: NO and NO2 are major part of a series of highly active primary and secondary compounds (including HCN and N2O5). Primary emission occurs mainly of NO which is rapidly converted to NO2, which thus becomes dominant in the atmosphere. Both these are quite short-lived species and are rapidly oxidized to nitrate aerosol or sulfuric acid. Both the gases are crucial in tropospheric and stratospheric ozone chemistry and in the chemistry of photochemical smog.
NO and NO2 are strongly influenced by anthropogenic emissions. Over 60% of nitrogen oxides come from combustion of fossil fuels and biomass. The amount of gases released from fossil-fuel combustion depends on the temperature of combustion process and nitrogen content of the fuel. Nitrogen content of coal is 1-2%, of crude oil is <1% and of natural gas is 5-10%. Concentrations of nitrogen oxides show high spatial variability during their short lifetime indicating that local and regional sources are highly important to their global budget. Natural sources of these oxides are soil and thermal dissociation of atmospheric N2 during lightening. Global emission of nitrogen oxides is about 50 Tg Y-1, which forms about 33% of total nitrogen, input into the atmosphere. About 43 Tg nitrogen is removed from atmosphere per year. This removal involves almost entirely the wet and dry deposition with a very small quantity lost to photochemical reactions. Concentrations of nitrogen oxides in clean ocean air in the troposphere are <100 pptv. Concentrations in rural air over the continents are 200-300 ppbv and in air influenced by human activities may be >10 ppbv reaching upto 500 ppbv in urban air. Highest concentrations are found in Northern Hemisphere around 400 N latitude where major anthropogenic sources of these oxides are located. Concentrations rapidly decrease with altitude to a background value of 10 pptv in the upper troposphere. Higher concentrations occur in winters, particularly in the mid-latitude areas under urban influence since temperature inversions are more prevalent and photochemical activity is at a minimum.
Table-4. Indicative characteristics of major atmospheric nitrogen compounds.
| Compound | Major sources | Nitrogen produced (Tg Y-1) | Background concentration | Polluted concentration | Lifetime | Sinks | |
| NH3 | Animals, soils, biomass burning | 54.0 | 0.1 ppbv | >6.0 ppbv | 6 days | Conversion to NH4 | |
| NH4+ | Conversion from NH3 | 65.0? | 0.05 g m-3 | >1.5 g m-3 | 5 days | Wet & dry deposition | |
| NO3- | Secondarily from NOx | 26.0 | 0.5 g m-3 | >10.0 g m-3 | 5 days | Wet & dry deposition | |
| N2O | Soil | 41.0 | 310 ppbv | - | 170 days | Strato-spheric photo-chemistry | |
| NO, NO2 | Fossil fuels, lightening, biomass burning, intercons-versions | 48.0 | <100 pptv | 100 pptv | <2 days | Oxidation to HNO3 & NO3-, photo-dissociation | |
Sulfur cycle
Most of the sulfur on Earth is stored in oceans (about 1.3 x 106 Pg), sedimentary rocks (about 2.7 x 106 Pg) and evaporites (about 5 x 106 Pg). Very small percentage reaches the surface and is exchanged with atmosphere. Accuracy of the natural emissions of sulfur is about 50% only.
Sulfur in lithosphere
Sulfur is 13th most abundant element in Earth’s crust (0.1%) and 9th most abundant in sediments. Sulfur content of rocks varies considerably e.g. sedimentary rocks have about 0.38% while igneous rocks have only 0.032%. Sulfur in lithosphere is mobilized by slow weathering of rock material. Dissolved in runoff, it moves with river-water and is deposited in continental shield sediments in oceans. Eventually on geological time-scale, this uplifts to surface again thus completing the geological part of the sulfur cycle.
Sulfur in hydrosphere
Main storage of sulfur in oceans is through dissolved sulfate, averaging about 2.7 g per kg. Most volatile sulfur compound in sea water is dimethyl sulfide (DMS; (CH3)2S) which is produced by algal and bacterial decay. Its concentration in sea water is about 100 x 10-9 L-1, highest concentrations being in coastal marshes and wetlands.
Sulfur is second most abundant compound in rivers with concentrations fluctuating highly with seasons and frequency of drought, flood and normal flow. Rivers transport about 100 Tg of sulfur per year to the oceans. The storage of main sulfur mass in oceans, sedimentary and evaporite rocks establishes the base for sulfur cycle.
Sulfur in soil and biosphere
Sulfur is major essential nutrient in the biosphere and is concentrated mainly in soil from where it enters biosphere through plant uptake. From soil, sulfur is also removed in solution to groundwater and by chemical volatilization. Its main sources are deposition from atmosphere, weathering of rocks, release from decay of organic matter and anthropogenic fertilizer, pesticides and irrigation water. In soil, it is present mainly in oxidized state (e.g. SO4-) with concentrations varying according to the amount of organic matter in soil. Rich organic soils may have upto 0.5% sulfur by dry weight.
Sulfur in soil may be in bound or unbound form, as organic or inorganic compounds, organic sulfur being most prevalent. Plants take up sulfur from the soil mainly as sulfate and it is passed on with the food chain in the biosphere. It leaves biosphere on death of living organisms when aerobic decay and decomposition brings back sulfate in the soil. Finally, anaerobic decomposition in soil releases part of organic sulfur as H2S, DMS and other organic compounds into the atmosphere. About 7 Tg of sulfur per year is released from global soils, with considerable latitudinal variation. The release of sulfur is dependent upon warmer temperatures.
Sulfur in atmosphere
Several sulfur compounds are released into the atmosphere due to interaction of processes between Earth’s surface and the atmosphere. Of these, most important six compounds are discussed below.
1. Carbonyl sulfide (COS): It is the most abundant sulfur species in atmosphere and in nature is mainly produced by decomposition processes in soil, marshes and wetlands along ocean coasts and areas of ocean upwelling that are rich in nutrients. Anthropogenic combustion processes produce less than 25% of COS. Its average concentration of about 500 pptv shows enough uniformity throughout latitudes and altitudes to suggest a long lifetime and no rapid sinks of this compound. A lifetime of 44 years is suggested with only sink being stratospheric photolysis and slow photochemical reactions in troposphere. Ocean may act both as source and sink. About 80% of total atmospheric sulfur is COS, but it is relatively inert and does not add much to atmospheric sulfur pollution problem.
2. Carbon disulfide (CS2 ):It is far more reactive than COS and has similar sources though on a smaller scale. It has lifetime of 12 days only and its major sink is photochemical reactions. As a result, CS2 shows greater spatial variation across the globe, ranging from 15 pptv in clean air to 190 pptv in polluted air. Its concentration decreases rapidly with altitude. The most important source of the compound is microbial processes in warm tropical soils. Major secondary sources are marshes and wetlands along sea coasts. Small anthropogenic inputs are from fossil fuel combustion.
3. Dimethyl sulfide (DMS): It is released from oceans in much greater amounts than COS or CS2 and has extremely small lifetime and is very rapidly oxidized to sulfur dioxide or is redeposited to oceans. In the sulfur cycle, most of natural gas released from oceans is DMS. Its concentrations are high during night, particularly in areas under some influence from continental sources.
4. Hydrogen sulfide (H2 S): It is mainly produced in nature during anaerobic decay in soils, wetlands, salt marshes and other areas of stagnant water with maximum concentrations occurring over tropical forests. This highly reactive is removed by reaction with hydroxyl radical (OH) and COS. Its highest concentrations occur at night and in early morning when photochemical activity is at a minimum.
4. Sulfur dioxide (SO2 ): Its natural source is oxidation of H2S and major anthropogenic source is combustion of fossil fuels. Its atmospheric concentrations are most influenced by anthropogenic emissions. In some industrialized areas such as eastern North America, over 90% of SO2 is from anthropogenic sources. Normally about half of global SO2 originates from natural sources. The lifetime of the gas is 2-4 days indicating that loss due to photochemical conversion to sulfate is quite important. Rest of the gas (about 45%) is removed from atmosphere by wet and dry deposition.
5. Sulfate aerosol: Sulfate aerosol particles originate from sea spray that is the largest natural source of sulfur to the atmosphere. Only 3 TG per year of sulfate is added to atmosphere from anthropogenic sources directly but much greater amounts are formed through secondary reactions from various sulfur species in atmosphere. Most of the salt spray sulfate falls back to oceans but some is carried over the continents to be included in deposition processes there.
Table-5. Indicative characteristics of major tropospheric sulfur compounds.
| Compound | Major sources | Sulfur produced (Tg Y-1) | Background concentration | Polluted concentration | Life-time | Sinks |
| COS | Soils, coastal marshes, biomass burning | 4.7 | 500 pptv | ? | 44 years | slow photoche-mistry, stratosphere, oceans |
| CS2 | Oceans, soils | 1.6 | 15-30 pptv | 100-200 pptv | 12 days | Photoche-mical production of SO2 |
| DMS | Oceans, algal deposition | 27-56 | <10 pptv | 100 pptv | 0.6 days | Oceans, oxidation to SO2 |
| H2S | Bacterial reduction, soils, wetlands | Variable | 30-100 pptv | 330-810 pptv | 4.4 days | Photoche-mistry |
| SO2 | Anthropo-genic sources, volcanoes, oxidation of H2S | 103 | 24-90 pptv | >5 ppbv | 2-4 days | Wet & dry deposition |
| SO4- | Sea-sprays, oxidation of SO2 | 138 | 0.1 g m-3 | >2.5 g m-3 | 1 week | Wet & dry deposition |
ORGANIC MATTER IN THE ENVIRONMENT
In the environment, organic matter is synthesized in its biotic component i.e. biosphere. Autotrophic organisms are the only organisms that can synthesize organic matter using solar radiation and mineral matter taken from atmosphere, hydrosphere and edaphosphere. Autotrophs synthesize organic matter either by photosynthesis or by chemosynthesis. While chemosynthesis is important for cycling of nitrogen and certain other processes in the environment, photosynthesis is the major process responsible for formation of organic matter in the environment. Autotrophic green plants, particularly land plants are most important from the point of view of photosynthetic production of organic matter. In the photosynthesis, carbon dioxide and water are used and a certain portion of short-wave solar radiation is absorbed and expended within the plant cover. In considering the role of plant cover of Earth in the global energy and water balance, it is necessary to consider the amount of solar radiation and water utilised by plants in production of biomass i.e. in photosynthesis. For this, the quantities that are calculated and studied are efficiency of photosynthesis and productivity of transpiration.
Efficiency of photosynthesis: It is the ratio of energy expenditure on the synthesis of biomass to the total quantity of solar energy absorbed by plant cover in an area. Many experimental studies have shown that this efficiency of photosynthesis is very modest and under normal conditions, usually does not exceed 0.1 to 1.0 percent. However, under very favorable conditions, it may increase to several percent.
Productivity of transpiration: It is the ratio of the amount of biomass produced to the quantity of water transpired by photosynthesizing plant cover. This productivity of transpiration usually ranges from 0.5 to 0.1 percent which indicates that photosynthesizing plants use very little water and abundant transpiration in them merely circulates the water in the environment.
Thus general low values of both the above quantities indicate that under natural conditions, plant cover assimilates only a negligible part of available energy and water resources i.e. there is substantial limitation on the use of natural resources in production of biomass in the environment. It is important to establish the causes of this limitation for the study of the relationship of productivity of plant cover to climatic factors. Experimental studies of maximum possible efficiency of photosynthesis in controlled environmental conditions when carbon dioxide of the atmosphere is fully utilized indicate that under such conditions, plants can assimilate 5% or more of the solar energy received and the productivity of transpiration also increases manifold. However, in natural conditions maximum possible photosynthesis and, therefore, the production of biomass is greatly limited by various factors other than the availability of resources.
PHOTOSYNTHESIS WITHIN TERRESTRIAL PLANT COVER
In nature, most of the photosynthesis takes place within the terrestrial plant cover in which different meteorological conditions exist at different levels. The efficiencies of photosynthesis at various levels within the plant cover are not same and are determined by particular microclimatic (meteorological) conditions prevailing at different levels. The microclimatic effects of a forest cover are explained in terms of:
Plant coverage characteristics: These characteristics depend upon:
Density of dominant forms in the forest covers.
Distribution of different forms in the forest covers.
Stratification characteristics of plant cover: These characteristics depend upon:
Total vertical height of plant cover.
Number of vertical strata in the plant cover.
Morphological characteristics of each strata in the plant cover which are determined by branching pattern of plants, evergreen or deciduous nature of foliage, size, density, texture and orientation of leaves.
Importance of the above features can be judged from comparison of tropical and temperate forest plant covers. In tropical forests, average height of tall trees is 46-55 metres, species diversity is 40-100 species per hectare, stratification is strong with 4-5 strata, undergrowth is dense commonly with two upper foliage strata and lower strata being denser. In temperate forests, average height of tall trees is about 30 metres, species diversity is less than 20 species per hectare, stratification is poor with usually 2-3 strata which are almost continuous from low shrubs to top of trees.
In the study of photosynthesis at various levels within plant cover, averaging of the values of meteorological elements at one level along the horizontal line is appropriate and it makes it possible to exclude the influence of individual plants on the meteorological regime. By applying such averaging techniques, following conclusions have been established:
Microclimate within the plant cover may be represented by a series of vertically varying profiles of meteorological elements, particularly of solar radiation, water vapor pressure, air temperature, carbon dioxide concentration and wind speed.
The profiles of meteorological elements show diurnal and seasonal variations.
The average vertical flow of short wave and long-wave radiation, heat and water vapor within the layer of plant covers and the momentum of the system depend substantially on height.
Microclimatic profiles within plant cover
Solar radiation: Plant cover significantly changes the pattern of incoming and outgoing radiation. Short-wave reflectivity of area depends somewhat on the density and characteristics of the plant cover. The albedo of areas having coniferous forests is about 8-14 while that having deciduous forests is about 12-18. Albedo of semiarid savannas and woodlands is much higher.
Large amount of solar radiation is trapped within the foliage canopy e.g. Fagus sylvetica forest traps about 80% of incoming radiation in the top strata of canopy and less than 5% reaches the ground. Such trapping is more pronounced on sunny days.
The foliage canopy absorbs more short-wave radiation than long-wave infrared radiation e.g. in tropical forests of Nigeria, only 7.6% radiation of <0.5 m reaches the ground while 45.3% of radiation of >0.6 m reaches the forest floor.
Effect of the age of plant cover on the penetration of light into the plant cover can be judged by the observation that in Pinus sylvestris forest in Germany, percentage of light reaching the ground floor is about 50% at 1.3 year, only 7.0% at 20 years and again 35% at 130 years.
Penetration of solar radiation within the plant cover generally obeys Bougner-Lambert Law:
I = Io e-KL
Where, I = radiation intensity on a horizontal plane within the plant cover; Io = radiation intensity on a horizontal plane above the plant cover; L = leaf area index; K = Extinction coefficient.
Extinction coefficient (K) is constant for a given species and is related to:
Amount and type of leaf chlorophyll.
Canopy architecture and
Reflectivity of leaves. Its value lies between 0.3 and 0.5 for grass-type plant cover and approaches 1.0 for nearly horizontal leaves. Value of K shows inverse relationship to chlorophyll content and reflectivity of leaves.
In general, light penetration into plant cover depends upon the type of plants (particularly trees), spacing of plants, age of plants, crown density, height of plants (particularly trees) and time of year. Percent light reaching the forest floor in some types of forests is given below:
Birch-beech forest 50-75%
Pine forest 20-40%
Spruce-fir forest 10-25%
Tropical forest 0.1-0.01%
In deciduous forests, light penetration increases during leafless conditions.
Thus the intensity of solar radiation decreases exponentially from top of plant cover towards Earth’s surface due to absorption and radiation scattering by the surface of plants. The resulting radiation balance, therefore, also decreases in the same direction due to screening effect of plants.
Air temperature: During day time, heating of foliage canopy causes a convectional transfer of sensible heat and so air temperature within upper canopy may be higher than above the canopy or below. At night, the relationship is reversed as upper canopy layer of air is cooled by contact which are both losing heat by radiation and also transpiring slowly.
Modification of thermal environment is due to shelter from sun, blanketing at night, heat loss by evotranspiration, reduction in wind speed and obstruction to vertical airflow.
Blanketing causes lower maximum and higher minimum temperatures and causes lower mean monthly temperatures in tropical and temperate forests.
At sea level, mean monthly differences in air temperature in temperate forest may reach 2.2OC in summer but only 0.1OC in winters. In hot summers, this difference can be more than 2.8OC.
In forests, which do not transpire greatly in summers e.g. forteto oak maquis of Mediterranean area, day temperatures in woods may cause mean monthly temperatures to be higher than in open.
Altitude in the same climatic zone may affect the degree of temperature decrease in temperate forests. At 1000 meters altitude, lowering of temperatures may be twice that at sea level.
Vertical stratification in plant cover modifies the thermal profile within it in complex ways. In tropical forests, dense foliage canopy heats up greatly during daytime and cools rapidly during night. It shows a much greater diurnal temperature range in denser canopy than in the lower strata. Whereas daily temperatures of second story are intermediate between those of the tree tops and undergrowth, the nocturnal minima are higher than either tree tops or undergrowth because the second story is insulated by trapped air both below and above.
Saturation water vapour pressure: This profile within plant cover shows close correspondence with temperature profile both during day and night. Forest temperatures differ strikingly from those in open and the forest water vapour pressures were found to be higher within an oak stand than outside it for every month except December.
At night, actual water vapor pressure almost reaches saturation as the air and canopies are cooled by radiation and convection. Some water vapor is transferred through transpiration from the canopy. During day, upper canopy is air heated by convection and water vapor pressure curve shifts much from saturation curve. Deficit between the two increases the downward and at quite lower level, actual water vapor curve inflects. Towards the bottom of canopy, it reapproaches saturation curve due to transpiration coupled with low air movement and low temperature towards base of plant cover.
The flow of water vapor within the plant cover increases with height because of the influence of transpiration by plants and the momentum of the system declines downwards from the plant cover’s upper boundary as a result of the inhibiting effect of plants on the movement of air. This effect is associated with the reduction in turbulent exchange within the layer of plant cover compared with higher air layers. The coefficient of turbulent exchange within the layer also declines towards Earth’s surface.
Humidity conditions within the plant cover are very much different from those outside it due to evotranspiration characteristics of the cover. It generally depends upon the type of plant cover, density of plant cover, structure of vertical stratification and temperature effects. Time of the day and season also affect evotranspiration and, therefore, humidity within the plant cover.
Evotranspiration generally increases with density of vegetation and within the plant cover, relative humidity may be 3-10% higher than outside. This effect is more pronounced in summers.
Rainforests have high transpiration and so have high humidity inside their plant cover. Mean annual relative humidity excess is reported to be 9.4% in beech, 8.6% in Pinus abies forest, 7.9% in larch forest and 3.9% in Pinus sylvestris forests.
In tropical forests, night exhibits complete saturation while in daytime, the humidity decreases with height.
Carbon dioxide: The profile of carbon dioxide concentration within plant cover shows much diurnal variation due to photosynthetic uptake of carbon dioxide during daytime and respiratory addition of this gas during night. Carbon dioxide concentration in soil is very low and its use by plants is spatially and temporally very inhomogeneous.
During daytime, CO2 concentration decreases from upper canopy towards ground. It reaches a minimum point near middle of canopy. Below this point, CO2 concentration rapidly increases towards ground and becomes equal to CO2 concentration outside the canopy at a point that roughly corresponds to compensation light intensity. It reaches fairly high level at soil surface. This profile is due to photosynthetic depletion of carbon dioxide in upper canopy, equilibrium corresponding to compensation point lower in the canopy and respiratory addition of carbon dioxide from lowest shaded leaves and soil microorganisms.
In the night, concentration of CO2 gradually increases towards ground level due to its respiratory addition.
Wind velocity: The profile of wind speed shows no strong change in day and night but overall wind speed is higher during daytime due to convectional effects. Wind profile within the canopy develops due to steady state boundary layer flow. The profile is logarithmic above the canopy and becomes exponential within the canopy. The zero plane displacement (D) depends on the height of plants. The roughness height (zo) is a measure of community roughness and it is effectively the thickness of a laminar sublayer through which individual elements project. Value of zo is related to height variation and spacing of individual elements which in the plant cover are plants. In extrapolation of logarithmic curve downwards, the zero velocity intercept is found to lie at the height D+zo. If the canopy were rigid, it would have a constant value but variation of surface roughness depends on leaf flutter, movement of branches and leaf streamlining. These variations cause variations in value of zo with wind speed. Surface frictional characteristics are entirely specified by D and zo. The wind profile is major factor in establishment of profiles of saturation water pressure and carbon dioxide within the plant cover.
Lateral air movement is generally lesser within the plant cover than outside it. Even large variations in outside wind velocities do not affect airflow inside forest cover. Vertical stratification structure, leaf canopy architecture, density of stand and season have marked influence on wind velocity profile within a plant cover. For this reason, reduction in wind velocity within the forest cover is different in temperate and tropical forests. Reduction in wind speed from outer edge towards deep inside a forest is greater in tropical rainforests. In temperate European forests, wind velocity at outer edge of forest is reduced to 60-80%, 50% and 7% at points 30 m, 60 m and 120 m respectively deep inside the forest. In Brazilian evergreen forest, wind velocity of 2.2 m/second at the outer edge of forest is reduced to 0.5 m/second at 100 m deep inside forest while at 1000 m inside forest the wind velocity becomes negligible. In this forest, outside storm velocity of 28 m/second was reduced to 2 m/second at 11 km deep inside the forest.
Flow of water vapor and momentum of system: The flow of water vapor within the layer of plant cover increases with height because of the influence of transpiration by plants. The momentum of the system decreases downwards from upper boundary of plant cover towards ground level as a result of the inhibiting effect of plants on the movement of air. This effect is associated with the reduction in the turbulent exchange within the layer of plant cover compared with higher air layers. The coefficient of turbulent exchange within the layer also decreases towards ground level.
The theory of photosynthesis within plant cover and numerical models of this process developed in recent years are based on the general idea of a transition from photosynthesis within a single leaf to photosynthesis within a layer that is homogeneous horizontally but possesses different physical conditions at various heights. Application of the theory of photosynthesis within a layer of plant cover indicates following general conclusions:
When assimilation process is not very sensitive to different meteorological elements, total assimilation within the plant cover strongly depends on the radiation flux for low levels of radiation. For large values of radiation, the total assimilation is independent of radiation flux and becomes dependent on other factors particularly temperature.
Within he plant cover, increase in total assimilation with increase in inflow of CO2 from soil is much slower than would occur if all the inflow of CO2 from soil were to be expended on assimilation. This is because the inflow of CO2 from soil first encounters the leaves located in shade, which are not able to photosynthesize intensively due to insufficient radiation. The general increase in CO2 concentration produced by its inflow from below is compensated by a reduced contribution of CO2 from above. Thus assimilation within plant cover is influenced very little by the upward inflow of CO2 from soil and is largely influenced by flow of CO2 coming downwards from the atmosphere.
PRODUCTIVITY OF PLANT COVER
The productivity of plant cover () is the difference between total assimilation and the expenditure of organic matter on respiration within a plant cover. Thus the productivity of a particular plant cover depends on photosynthesis and respiration in it. The leaves of plants are the major organs association with both photosynthesis and respiration. Therefore, the productivity of plant cover substantially depends on the value of the index of leaf surface (leaf index) and decreases for both very small and very large values of this index. In view of this, the value of productivity of a plant cover is calculated for an optimal value of leaf index i.e. the value of this index that corresponds to the highest value of productivity. From various studies, it has been established that the parameters and factors that affect the photosynthesis within a plant cover also influence the productivity of plant cover. Thus the productivity of plant cover is mainly determined by parameters characterizing the properties of plant cover itself and the climate.
In general, following important points can be observed in relation to productivity of plant cover in nature:
The structure of plants in the plant cover continuously changes throughout their life cycles and photosynthetic activity of leaves is never optimal throughout the entire vegetative period of any plant.
Availability of mineral nutrients in nature is always less than optimally required for maximum possible photosynthesis.
Under natural conditions, water regime of soil is also not constantly maintained at optimally required level.
Thus the productivity of plant cover in real natural conditions is always less than theoretically possible maximum level due to complex interactions between a variety of biological, climatic and soil factors.
Climatic factors and productivity of plant cover
In conditions of sufficient moisture, two climatic factors i.e. photosynthetically active radiation and temperature are particularly important in relation to productivity of plant cover.
The influence of radiation and temperature on productivity of plant cover is quite complex. In real natural situations, radiation is always a factor whose value is a ‘minimum’ because radiation available to leaves in lower layers of canopy is always insufficient. Therefore, increase in radiation flux always results in increased productivity of plant cover.
With increase in temperature, the productivity of plant cover increases initially. After attaining a certain maximum value that depends on the value of radiation flux, productivity begins to decrease with further increase in temperature. Thus productivity of plant cover substantially decreases above a certain threshold value of temperature which is determined by the radiation flux.
PRODUCTIVITY OF GLOBAL PLANT COVERS
Average values of the productivity of natural plant covers of Earth have been derived by using various theoretical and numerical models and data from a variety of studies including empirical determinations of productivity in individual biogeographical zones.
Terrestrial plant covers
Yefimova (1979) has made use of quite precise relationships between productivity of natural plant cover and meteorological factors in calculating values of the productivity and the coefficient of utilisation of photosynthetically active radiation for each continent. Results of her calculations are given in the Table-6. The data shows that the average productivity per unit area for the five continents of Earth does not differ very much. In each of these continents, magnitude of productivity over large part of continental territory is greatly limited by insufficient moisture or heat. The continent of South America is exception to this general condition since climatic conditions over large part of its territory are favourable for plant life.
Table 6: Productivity and coefficients of utilization of photosynthetically active radiation in various continents of Earth. (Yefimova 1979)
Continent | Productivity (x109 tonnes) | Productivity (center per hectare) | Coefficient of utilization of photosynthetically active radiation (as %age of total over vegetative period) |
| Europe Asia Africa North America South America Australia (including islands of Oceania) | 8.9 38.3 31.0 18.1 37.2 7.6 | 85 98 103 82 209 86 | 1.26 0.88 0.59 0.94 1.13 0.44
|
In Australia and Africa, coefficients of utilization of photosynthetically active radiation are lower than average. This can be attributed to insufficient moisture over large parts of these continents, which inhibits the complete utilization of available radiation by plant covers.
In Europe and South America, most favorable conditions for the development of plant life are found. In Europe, located at higher latitudes and exposed to less solar radiation, its utilization is relatively greater.
Smil (1985) gave estimates of the productivity and storage of biomass in major biomes of the Earth. These estimates are given in Table-7. Data in this table shows that there is not much difference in the area occupied by different types of ecosystems except wetlands that occupy smallest area on the Earth. However, productivity is highest in cultivated lands where one ton of biomass is produced per one ton of phytomass, followed by tropical and temperate grasslands where 0.5 ton of biomass is produced by each ton of phytomass. Next in productivity are tundra, deserts-semi-deserts and wetlands where 0.2 tonnes of biomass is produced per hectare from one ton of phytomass per hectare. These areas are followed by wetlands and shrub-lands where productivity is 0.13 tonnes per hectare. Tropical, temperate and boreal forest, though occupy almost same area on Earth, produce 0.067, 0.04 and 0.02 ton of biomass per tone of phytomass per hectare respectively. Despite these facts, most important on Earth are tropical, temperate and boreal forests that have the highest concentration of biomass on Earth (totaling about 750 tonnes per hectare). These ecosystems also have the highest total storage of biomass on Earth totaling about 850 x 109 tonnes. Further, it may be noted that contribution to total biomass production is equal for tropical rainforests and tropical grasslands (20 x 109 t/yr), followed by boreal forests and tropical grasslands (15 x 109 t/yr) and temperate forests, woodlands-shrub-lands and temperate grasslands (10 x 109 t/yr). Tundra and deserts have quite high average of net biomass production per unit area and also quite high weight of phytomass per unit area. Despite this they contribute very little to total global biomass production (1.0 – 2.0 x109 t/yr). However, if total biomass storage in different types of ecosystems on Earth is considered, tropical rainforests, temperate forests and boreal forests are the most important storehouses of organic matter on Earth having 850x109 tonnes of biomass. Woodland and shrub-lands having 75x109 tonnes and then tropical and temperate grasslands having 60x109 tonnes of biomass storage follow these.
Table 7: Area, productivity and storage of major global ecosystems. (Smil, 1985)
| Ecosystem | Total area (x106 km2) | Average net production (tonnes/ha) | Average phytomass (tonnes/ha) | Total production (x109 tonnes/year) | Total storage (x109 tonnes/year) |
| Tropical rainforest
Temperate forests
Boreal forests
Woodland and shrub-land
Tropical grasslands
Temperate grasslands
Cultivation
Tundra
Deserts and semi-deserts
Wetlands
Settlements and transport
| 10.0
10.0
15.0
10.0
10.0
10.0
15.0
10.0
20.0
5.0
5.0
| 20.0
10.0
10.0
10.0
10.0
10.0
10.0
1.0
1.0
15.0
5.0
| 300.0
250.0
200.0
75.0
20.0
20.0
10.0
5.0
5.0
75.0
5.0 | 20.0
10.0
15.0
10.0
20.0
10.0
15.0
1.0
2.0
8.0
3.0 | 300.0
250.0
300.0
75.0
40.0
20.0
15.0
5.0
10.0
40.0
3.0
|
| Total |
|
|
| 114.0 | 1058.0 |
Aquatic plant covers
There is much less data about productivity of autotrophic plant covers in water bodies as compared to that about terrestrial plant covers. However, the available data indicates that the seas and oceans have the greatest volume of organic matter produced by phytoplankton located in the 30-40 meters deep layer of hydrosphere. At greater depths, quantity of solar radiation is insufficient for active development of photosynthesis.
In general, the productivity of shelf zones is substantially less than open ocean. It may attain maximum values in small bodies of water possessing large quantities of minerals required by the plants. The overall value of productivity for the oceans is estimated to be about 55 billion tonnes per year i.e. approximately 15 centner per hectare. This last figure is less than 1/6th of the average productivity per unit area on continents.
Thus the estimates show that the yearly volume of productivity for the Earth as a whole is approximately 200 billion tonnes i.e. about 40 calories per hectare. This corresponds to an energy expenditure of approximately 0.15 kcal/cm2 per year. This is about 0.1% of the solar radiation reaching the Earth’s surface.
TRANSFORMATIONS OF ORGANIC MATTER
In the ecosystem, autotrophic organisms (chiefly the green plants) use the energy of solar radiation to produce organic matter, which is used by all the living organisms, including autotrophic organisms themselves, in running their life activities. The organic matter is used by the living organisms through their respiratory activity. Out the total organic matter produced by photosynthetic activity of autotrophic organisms, a certain portion is consumed by these organisms themselves and the remaining organic matter is available in the ecosystem as the net organic matter production of autotrophic organisms. A relatively very small part of this net organic matter production in the ecosystem is directly transformed into mineral substances. This takes place without the participation of any other living organisms through the processes such as forest and prairie fires during which organic matter is transformed into carbon dioxide, water vapor and certain mineral compounds. Further, a still smaller portion of organic matter is deposited in the upper layers of lithosphere and at the bottom of water bodies in the form of coal, peat and other organic compounds. The remaining organic matter is now passed on to heterotrophic organisms in the ecosystem through various food chains. All the living organisms of a particular type in the ecosystem that receive organic matter as food in a particular manner constitute a trophic level. The organic matter received by a trophic level undergoes three fates:
A portion is consumed by that that trophic level itself though respiration in that trophic level
A certain other portion is passed on to next higher trophic level as organic food and
Remaining organic matter is stored in the trophic level as increase in the biomass of that trophic level (i.e. increase in the number of organisms of that trophic level).
From the point of view of ecosystem energetics, the organic matter that is received, passed on to next higher trophic level or stored by a trophic level represents the amount of energy received, passed on or stored by that trophic level. It is obvious that in a dynamically stable ecosystem, there can not be any storage of energy (i.e. organic matter) in any of its trophic levels. Therefore, in the dynamically stable global ecosystem, a very small portion of the net production of organic matter by autotrophs is stored in the abiotic components of the environment (i.e. lithosphere and hydrosphere) while major portion is consumed by heterotrophic organisms through their respiration.
The consumption of organic matter in a trophic level (including autotrophic organisms themselves) through the respiration in that level represents the loss of energy in that trophic level. It is a feature of global ecosystem that the flow of energy (represented by flow of organic matter as food) between trophic levels is associated with large losses of energy at each trophic level. The ratio of the amount of energy passed on from a trophic level to its next higher trophic level (n) and the amount of energy received by that trophic level from its previous trophic level (n1) is termed ecological efficiency () of that trophic level i.e.
Ecological efficiency () = n/n-1
The ecological efficiency of trophic levels, in general, is estimated to range between 10-20%. Such small general value of ecological efficiency indicates that biomass in each successively higher trophic level in the ecosystem is bound to be substantially reduced. Since ecological efficiency of a trophic level depends on the respiration of that level, smaller the value of ecological efficiency of a trophic level, greater is the consumption of organic matter through respiration (i.e. loss of energy) at that trophic level. As a result, there is greater reduction of biomass in that trophic level and in the next higher trophic level.
Nature of organisms and transformation of organic matter
Since intensity of metabolism per unit mass of a live organism usually increases with decrease in the size of organism, the biomass present at a specific trophic level in the food chain depends on the size of organisms of that trophic level. One of the causes of this relationship is that the metabolism depends substantially on the ratio of the rate of diffusion of gases through the surface and the mass of organism. This ratio increases as the size of organism decreases. Thus the rate of metabolism of a given unit weight of microorganisms is many times greater than that of macro-organisms. Further, metabolism also depends on the nature of physiological processes within the tissues of organisms. In wood of plants, the metabolism is usually much slower than in vertebrate tissue of similar size. These general principals largely determine the total biomass of various types of organisms in the global ecosystem.
The largest proportion of forests in the overall biomass of living organisms is due to the fact that autotrophic trees are located at the first link in the food chains and also due to the large size of individual trees. Together with specific properties of the wood, this feature substantially reduces the rate of metabolism per unit biomass in forests. Though the productivity of ocean phytoplankton is comparable with forests, small size of individual plankton organisms intensifies their metabolism per unit weight so much that the total mass of plankton on Earth is negligible in comparison with that of forests.
About 95% of the total biomass on Earth belongs to plants and rest to the animals. Biomass of aquatic organisms is substantially less than that of terrestrial organisms. Therefore, the distribution global biomass is largely determined by the distribution of terrestrial plant cover i.e. by the forest cover on continents. Considering that total biomass on Earth (global biomass) is approximately 3x1012 tonnes and total productivity of plants on continents is approximately 140x109 tonnes, the time period of one cycle of organic matter for the plants on Earth comes to be approximately 20 years. This average figure relates to forests that constitute major portion of the biomass of plants on Earth. In other natural zones on continents, the duration of one cycle of plant organic matter is much shorter. The duration of this cycle in the oceans having phytoplankton is still shorter and appears to be only a few days.
The total biomass of animals is assumed to be approximately 1011 tonnes. Assuming that the animals assimilate about 10% of the total productivity of plants, the average duration of one cycle of animal organic matter comes to be several years. However, the actual length of life of one generation varies widely in animal kingdom and the nature of the distribution of biomass among different animal groups is still not much clear.
Invertebrates are the largest components of animal biomass and among them, most important are organisms living in soil. The zoological mass of large animals per unit area on Earth is relatively quite low. Calculations of Huxley (1962) show that while in African savannahs, the biomass of large wild animals may be 15-25 tonnes/km2, this figure is only about 1.0 ton/km2 in middle latitudes, 0.8 ton/km2 in tundra and 0.35 ton/km2 in semi-desert areas.
Man occupies topmost position in the food chain on the Earth and consumes both the primary production of autotrophic plants and the biomass produced by many herbivorous and carnivorous animals. For the present size of human population of over 4.0 billion, its biomass is approximately 0.2x109 tonnes. Assuming that each human being expends on average about 2.5x103 kcal of energy per day, the total energy consumption of human population comes to be about 1.8x1015 kcal/year. Thus, the human population consumes about 0.2% of the total production of Earth’s organic world.
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