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The Natural History of Oxygen
M A L C O L M DOLE
From the Department of Chemistry and the Materials Research Center, Northwestern University, Evanston
ABSTRACT The nuclear reactions occurring in the cores of stars which are believed to produce the element oxygen are first described. Evidence for the absence of free oxygen in the early atmosphere of the earth is reviewed. Mech- anisms of creation of atmospheric oxygen by photochemical processes are then discussed in detail. Uncertainty regarding the rate of diffusion of water vapor through the cold trap at 70 km altitude in calculating the rate of the photo- chemical production of oxygen is avoided by using data for the concentration of hydrogen atoms at 90 km obtained from the Meinel O H absorption bands. It is estimated that the present atmospheric oxygen content could have been produced five to ten times during the earth's history. It is shown that the isotopic composition of atmospheric oxygen is not that of photosynthetic oxygen. The fractionation of oxygen isotopes by organic respiration and oxidation occurs in a direction to enhance the O is content of the atmosphere and compensates for the 018 dilution resulting from photosynthetic oxygen. Thus, an oxygen isotope cycle exists in nature.
I. I N T R O D U C T I O N
Although it is estimated that oxygen is only the third most a b u n d a n t element cosmically (1) coming behind hydrogen and helium, in that order, it is the most a b u n d a n t element on the earth's crust, which is important for the living race as oxygen is essential for life. As we shall see later, oxygen was very p r o b a b l y not present in the early atmosphere of the earth, so it is interesting to consider how oxygen arose in the atmosphere, whether its a b u n d a n c e is now changing, w h a t its isotopic composition is, and how the latter varies between samples of oxygen from different sources. Furthermore, there is p r o b a b l y a close connection between the origin of life on this planet and the growth in the a b u n d a n c e of atmos- pheric oxygen as recently emphasized b y Berkner and Marshall (2). Before con- sidering the interesting factors involved in the development of atmospheric oxygen, however, we shall take up first the history of the formation of the element oxygen in the cosmos. It should be emphasized that m a n y of the con- clusions described below must be considered not in the same light in which we view well grounded scientific laws and principles, b u t rather as the best
The Journal of General Physiology
STRUCTURE AND FUNCTION OF OXYGEN
intelligent guesses and deductions that we can now make on the basis of presently available facts and theory.
II. C O S M I C F O R M A T I O N O F O X Y G E N
It is believed that the energy radiated b y a star throughout most of its life comes from the so called hydrogen-burning nuclear reaction in which hy- drogen is converted to helium (3). If this is true, then at zero time the universe must have consisted largely and probably solely of hydrogen, because there are few nuclear reactions which spontaneously produce hydrogen. In other words, hydrogen is continually being converted to helium and heavier ele- ments and because it is the most a b u n d a n t element in the universe today it must have been present in overwhelming a m o u n t and very possibly completely pure at zero time. The hydrogen-burning reaction in which helium is produced is the most efficient energy producer, but the actual mechanism probably depends on the temperature of the star. At relatively low temperatures and in the early stages of a star's life before m u c h H e 4 has been produced, the hydrogen- burning chain reaction consists of the following steps:
1H 1 + 1H 1 - - o l H ~ + / 3 + + v
where/3 + represents the positron and v the neutrino,
1H2 + 1H 1 --~ ~He 3 + "y ( 2 )
and
~He ~ + 2He 3 ~ H e 4 + 2 1 H 1 ( 3 )
where 3' represents a g a m m a ray photon. The total energy yield is 26.2 M e v (equivalent to 6 X l0 s kcal per mole of 2He 4) for the net process which is
41H J --o 2He 4 + 2~3+ + 2v + 23/ ( 4 )
(There is a 2 per cent energy loss due to the production of the neutrino.) There are other mechanisms for H e formation but the above is believed to be the most important. As the hydrogen is consumed in the core of the star to form helium, no further nuclear transformations can take place until both the temperature and density have greatly increased. At temperatures of 108°K and densities of 105 gm cm -3 sC 12 can be produced by the nuclear reaction
32He4 ~ ~Ct2* --o 6C'2 + T (5)
with an energy release of 7.3 M e v per atom of ~C'2. This reaction probably
S T R U G T U R E A N D F U N C T I O N O F O X Y G E N
c a n be d e t e c t e d , therefore, t h r o u g h n u c l e a r m a g n e t i c r e s o n a n c e m e a s u r e - ments. W e n o w pass t h r o u g h t h e speculative phases (7) in the e a r t h ' s h i s t o r y of its f o r m a t i o n b e g i n n i n g w i t h the b r e a k i n g u p of the solar n e b u l a , c o n t i n u i n g o n to the c o n d e n s a t i o n of some of its m a t t e r into p r o t o p l a n e t s , the slow e v a p o - r a t i o n a n d escape of 99 p e r c e n t of t h e mass o f the e a r t h ' s p r o t o p l a n e t into i n t e r p l a n e t a r y space lasting a b o u t 10, years until we c o m e to the e a r t h with its p r e s e n t mass a n d p r i m i t i v e a t m o s p h e r e.
III. THE EARLY ATMOSPHERE OF THE EARTH
The earth now has an atmosphere completely different from that of any other celestial body in the universe whose atmosphere is known (see Table II), but
T A B L E I I COMPOSITION OF PLANET ATMOSPHERES (8) Units, centimeters of gas at standard temperature and p r e s s u r e above the visible surface* Mean distance Planet from sun H~ Ar N2 02 C O , CH4 N H , H 2 0 miles X 10-t Mercury 35. Venus 67.24? <200 50,000 < Earth 92.91 <1 7,400 625,000 168,000 200 2 Mars 141.74 2,000 178,000 <200 420 < Jupiter 483.90 D 15, Saturn 887.2 D 35, Uranus 1,785.0 2X l0 s 150, Neptune 2,797.0 D 250, Pluto 3,670.
~-,<4 <I
<2 i < <
- D signifies that the gas is undetectable but probably is a dominating component.
it was n o t always so u n i q u e. T h e r e is a b u n d a n t e v i d e n c e t h a t o x y g e n was n o t p r e s e n t in the original e a r t h ' s a t m o s p h e r e. C h a m b e r l i n (9) has s u m m a r i z e d t h e facts w h i c h briefly m a y b e listed as follows: t h e e a r t h ' s crust is still sub- oxidized, the rocks of w h i c h c a n a n d d o take u p o x y g e n p r i m a r i l y to oxidize ferrous i r o n to ferric a n d to a lesser e x t e n t to oxidize m a n g a n o u s a n d sulfur c o m p o u n d s. F e r r o u s iron in h o t l a v a absorbs free o x y g e n a n d e v e n extracts o x y g e n f r o m w a t e r to f o r m Fe304 a n d H2. T h e r e a c t i o n s in a h o t v o l c a n i c e r u p t i o n are a n i m p o r t a n t source of free h y d r o g e n in the a t m o s p h e r e. T h e o x y g e n t h a t t o d a y is f o u n d in the a t m o s p h e r e was p r o d u c e d b y the p h o t o - s y n t h e t i c r e a c t i o n
CO, nt- H 2 0 -t- hu chlorophyll (^) , O~ + - (^1) (CH20). n
MALCOLM DOLE NaturalHistory of Oxygen 9
in which 1 mole of the product oxygen is chemically equivalent to 1 mole of the reactant carbon dioxide. If the carbohydrates formed in the reaction eventually are reduced to coal or fossil fuels, then the sum of the carbon in the latter should be chemically equivalent to the free oxygen of the atmos- phere. Actually a considerable excess of carbon is now recognized, hence there is no necessity for assuming the presence of any free oxygen in the initial atmosphere of the earth. In fact, Rubey (10) has estimated that the entire carbon-oxygen ratio of the earth's crust, atmosphere, hydro-, and biospheres, can be accounted for by an approximately 3 to 2 ratio of CO2 and CO "which is within the range of occluded CO2 and CO proportions actually found in igneous rocks." The reason that we made the statement above that all of the oxygen of the present atmosphere was produced by the photosynthetic reaction is that only approximately 2000 years are required to produce all of the atmospheric oxygen (2). This is a very short time on the geologic time scale. Two other arguments against the existence of free atmospheric oxygen in the primeval atmosphere are the following: It is generally assumed that life on this planet could not have arisen in an oxidizing atmosphere. For example, Miller (11) was able to synthesize amino acids by passing an electric spark through a mixture of CH4, NH3, H2, and H~O, but no organic compounds were formed when the gas mixture was only CO2 and H20. Abelson (12) found that CH4 -b N2 + H20 gave amino acids in an electric discharge, but not when COs -I- N~ -b H20 was the gas mixture. In the presence of oxygen and light, especially ultraviolet light, the lifetime of amino acids would be very short. Thus, these authors believe that for life to have started, free oxygen must have been absent from the early atmosphere. The second reason for the absence of oxygen in the early atmosphere stems from the deduction of Brown (13) that because most of the neon and a large fraction of the other inert gases escaped from the earth during its formation from its protoplanet, no free gas could have existed in the atmosphere. Only those gases were retained that could be retained chemically, H20, CO2, 02, and N2. For oxygen to exist as free gaseous oxygen it had to be liberated from chemical combination. Possible mechanisms for such liberation via photo- chemical means are discussed in the next section.
IV. T H E P H O T O C H E M I C A L T H E O R Y O F T H E F O R M A T I O N O F A T M O S P H E R I C O X Y G E N A. Photochemical Reactions Although the greatest energy of the sunlight striking the earth per 50 A range of wavelength is in the visible (2) at about 4500 A wavelength, nevertheless, there is considerable energy in the ultraviolet region of the spectrum. As the wavelength is decreased to 2000 A or lower, water and then carbon dioxide
MALCOLM DOLE Natural History of Oxygen z I
Further reaction of these radicals would lead to water formation. T h e yield of the reaction is one-half molecule of hydrogen reacted per q u a n t u m of light absorbed by the ozone (15). For each molecule of water formed one atom of free oxygen is lost and reactions (11) and (12) essentially reversed. Obviously, for oxygen to build up in the atmosphere it is necessary for hydrogen to escape into outer space. T h e author some time ago (14) con- sidered m a n y possible reactions that might trap atomic hydrogen and pre-
,
=E o I 0
1,.-, z __. tl.
o^ o
z^ i O - I o I -
o m~ 10-
1 0 0 3
I 0 0 0
4500 102
H
o ~ A,o
[ A T o
I (^) 1200 I (^) I (^) 1400 I (^) I (^) 1600 I (^) I (^) 1800 I (^)! (^) 2 0 0 0 I (^) I " " ~ , 2 2 0 0 I WAVELENGTH ( A } Floum~ l. Ultraviolet absorption coefficientsat different wavelengths for some atmo-
spheric gases from Berkner and Marshall (2). Figure reprinted by permission of the Faraday
Society from Discussions o/the Faraday Society, 1964, 37, 122.
vent it from escaping, but he could find none that looked very plausible. Even at heights in the atmosphere where both ozone (16) and atomic hydrogen can be produced, about 90 km, the following reactions are believed to occur (17)
Oa + H --~ O2 + O H --3. OH + O ~ O 2 + H --0.
T h e net result is simply Os + O --* 2 O2 leaving the H a t o m concentration unchanged. Evidence for these reactions is found in the Meinel O H bands at
I 2 S T R U C T U R E A N D P U N C T I O N O F O X Y G E N
10,400 A in the spectrum of the night sky (17). The energy, 3.30 ev of reaction (21), is just sufficient to excite O H to the 9th vibrational level and the transi- tions from this level are found to be the most intense.
B. Rate of Hydrogen Escape The velocity that a n y body must have to escape from the earth, v e, is given by the equation (18)
ve = "~/2 gro (23)
where g is the acceleration due to gravity, 981 cm sec. -2, and ro is the m e a n radius of the earth, 6.38 × l0 s cm. The escape velocity is calculated to be 11.2 km sec. -1 or about 24,000 miles per hour. T h e temperature required so that the greatest number of hydrogen atoms would have this velocity is 7500°K, an enormously high temperature. But due to the Boltzmann distri- bution even at normal temperatures a certain fraction of molecules will have velocities equal to the escape velocity. The equation for the n u m b e r of mole- cules leaving unit area per second with an upward component of velocity and speeds greater than ve is (18)
n (/~2ve2-b 1)e -~2v"2 ( 2 4 )
where n is the number of atoms per cm 3 at a height in the atmosphere where the temperature is equal to T ° K and /3 is defined by the equation
/ 3 ~ M (25)
2 R T
where M is the molecular weight of the gas and R is the gas constant. At a height of 90 km in the atmosphere where the pressure is about 10 - a t m and the temperature about 200°K, Harteck and Reeves (19) have esti- mated that there are 1.5 × 109 H atoms cm -3. Under these conditions the number of H atoms escaping per cm 2 per second is 0.15. At this height a shell about the earth would have a surface area of 5.26 X 1018 cm 2 so that the total number of H atoms escaping per second would be 7.50 X 1017 or in terms of moles o f H , O , 0.62 X 10 -6. There are 3.15 X 107 seconds per year, so that 19.6 moles of water are lost per year. Inasmuch as there are 7.6 X 10 ~ gm atoms of oxygen in the present atmosphere, it would take 7.6 × 1019/19. or about 4 X 1018 years to produce all of the oxygen of the atmosphere by this mechanism, a fantastically long time. However, if the hydrogen atoms are at a height of about 125 km where the temperature is 500°K (20), then the number escaping per cm ~ per second becomes 5.8 X 108 H atoms, or
I4 STRUCTURE^ AND^ FUNCTION^ OF^ OXYGEN
mate of the water vapor pressure considerably lower than that of Harteck or Kuiper, he deduced a rate of H a t o m escape a b o u t 500 times less than the rates given above. However, any turbulence or mixing in the stratosphere would raise Urey's estimate. But the whole problem of the partial pressure of water vapor can be avoided b y considering only the rate of H a t o m escape as explained below. Although some of the H atoms of reaction (2 I) m a y have come from the photodissociation of CH4, the low a b u n d a n c e of the latter and the probability of its reaction with ozone or nitrogen oxides at lower levels indicate that this source of H atoms can be neglected. But for the 0 2 photochemical process to leave a residue of free atmospheric oxygen, the H atoms must diffuse from the lower cold levels to higher hot levels where their kinetic energy will be sufficient for escape. Using equation
T A B L E I I I R E S U L T S O F E S C A P E A N D D I F F U S I O N C A L C U L A T I O N S (Assuming H a t o m c o n c e n t r a t i o n at 90 k m to be 1.5 X I0 ~ a t o m s / c m 8) H a t o m escape Conditions (^) No. of H atoms Yrs. to reproduce O~ in Height Temperature escaping/cmS sec. present atmosphere
90 k m 200°K 0.15 4X 125 500 5. 8 0 X 108 0. 9 5 X 109 W a t e r diffusion c a l c u l a t i o n s 20 260 ( H a r t e c k a n d 0. 8 3 X 108 J e n s e n , 1948) 40 270 ( K u i p e r , 1952) 0.62X H a t o m diffusion 90 220 (average) 1.5XlO 9 0. 4 2 X 109
E s t i m a t e d age of e a r t h 5 X l O 9
(26) to make this calculation from the 90 k m level assuming the H atom concentration to be 1.5 × 10Vcm ~ as before, it turns out that 0.42 × 109 years would be required for enough H atoms to diffuse so that the present atmospheric O2 content could be generated. T h e results of the diffusion and escape calculations are collected in Table III. If the H atom concentration is 1.5 × 109 a t o m s / c m 3 at 90 km, the con- centration would probably be somewhat less than this at 125 km; therefore, the calculation of the rate of the H atom escape at 125 km m a y be too fast b y the ratio of these concentrations. Nevertheless, it would appear that these new estimates based on recent determinations of the H atom a b u n d a n c e at 90 km confirm the earlier deductions of Harteck and Jensen and of Kuiper, that the photochemical dissociation of water vapor to oxygen and hydrogen during the lifetime of the earth has been of considerable geological signifi- cance.
MALCOLM DOLE Natural History of Oxygen 15
V. O X Y G E N A N D T H E O R I G I N O F L I F E
A. Obstades to Life in the Earliest Times
Intelligent life as we know it requires an extremely specialized environment whose temperature and composition have to be carefully regulated within rather narrow limits for life to survive, as was pointed out in a masterly fashion m a n y years ago b y L. J. Henderson (22). In the earliest times be- cause of less intense radiation from the sun, the earth m a y have been somewhat colder than it is today (23), in which case the oceans m a y have either not
yet been formed or m a y have been too cold, i.e. frozen, so that photosynthesis
could not have occurred. T h e initiation of life m a y have had to await, there- fore, a certain warming up of the earth. Another obstacle to the creation of life in the earliest times must have been the intense ultraviolet radiation striking the earth's surface. In Fig. 1 the absorption coefficients in the ultraviolet are plotted for a n u m b e r of gases and it can be seen that only ozone has a significant absorption coefficient for wavelengths greater than a b o u t 2000 A. In fact, the absorption coefficient of ozone goes through a m a x i m u m (24) at 2537 A equal to a b o u t 135 c m -1. This means that in the early atmosphere before the high altitude accumulation of ozone, the ultraviolet intensity at the earth's surface must have been great enough to be lethal to m a n y organisms. Especially the combination of oxygen, even ff the latter is present only at low concentrations, with ultraviolet light would be particularly deleterious (12). Berkner and Marshall (2) have re- cently reemphasized the need for a protective ozone layer to shield the earth from the lethal ultraviolet radiation. M o r e importantly, they have shown a possible connection between the various stages in the evolution of life and the levels of oxygen concentration in the atmosphere. It should be noted, how- ever, that U r e y (21) has suggested that there m a y have been organic com- ponents in the early atmosphere of sufficient a b u n d a n c e to have provided a primeval ultraviolet screen.
B. The Early Chemistry of Atmospheric Oxygen
As soon as the water vapor content of the atmosphere rose to significant levels,
i.e. as soon as the earth w a r m e d up if it had been cold, the photochemical
process for the production of oxygen must have started. At altitudes of 10 k m and above the processes going on would be little affected if at all b y the con- dition of the earth's surface. As soon as oxygen was produced it would begin diffusing downwards. As soon as it reached the top of the troposphere, a b o u t 10 kin, it would rapidly mix with all the gases of the troposphere because of winds and turbulence in the latter. Because the formation of free atmospheric oxygen depends on the escape of hydrogen from the earth and because the latter m a y have occurred slowly,
MALCOLM DOLE NaturalHistory of Oxygen 17
any chemical reaction, that the element oxygen always had exactly the same
ratio of isotopes irrespective of the substance in which the oxygen was com-
bined. Even after the discovery of heavy hydrogen by Urey, Brickwedde, and
M u r p h y (25) and of the rather rapid electrolysis separation method by Wash-
burn and Urey (26) which was carried to an exciting completion by G. N.
Lewis (27), there was still some question whether such separations would be
possible with other elements. T h e author recalls talking with the late Pro-
fessor R. H. Fowler of the University of Cambridge in the summer of 1933,
just after Fowler had returned there from a visit to Lewis at Berkeley. Fowler
suggested then that whereas light and heavy hydrogen could be separated
because of the 100 per cent differences in their masses, this might not be pos-
sible with any other elements. In other words, the hydrogen isotopes might
be unique.
T A B L E I V ISOTOPIC C O M P O S I T I O N OF OXYGEN FROM VARIOUS SOURCES Source Atomic weights A t o m per cent 01s
Air Photosynthesis (in Lake Michigan water) Lake Michigan water Atlantic Ocean water Carbonate rocks CO2 (at 0°C in the presence of fresh water)
15.99938(6) 0. 15.99928 0. 15.99926 0. 15.99928 0. 15.9994o 0. 15.9994~ 0.
However, Urey and Greiff (28) computed the partition function ratios of
various isotope pairs such as [018]/[O16] and [H~O181/[geo 16] and concluded
that in m a n y isotopic exchange equilibrium reactions like the following
2I-I~OlS(g) +^ 0 1 6 z " , ~gj ~- 214~o~(g)^ x^ + o~ (g)^18 (28)
the equilibrium constant would be slightly greater than unity. Within a
m o n t h or two after the publication of this paper the author (29, 30) dis-
covered a significant difference in the atomic weight of oxygen in the air as
compared to the oxygen in Lake Michigan water. A review contains more
recent data (31). Other sources of oxygen such as the oxygen in carbonate
rocks and iron oxide ores gave still different oxygen isotope ratios (32). Obvi-
ously, the element oxygen was not a good standard for the atomic weight
scale so the International Committee on Atomic Weights in 1961 abandoned
(6) oxygen and chose the carbon-12 isotope as the standard, giving to it an
atomic weight of exactly 12.0000.
Table I V summarizes the atomic weights of oxygen from various sources.
i8 S T R U C T U R E A N D F U N C T I O N O F O X Y G E N
The International Committee chose air as the source of oxygen for their
table. Dole, Lane, Rudd, and Zaukelies (33) showed that oxygen in air
samples taken from all over the world and up to altitudes of 87,000 feet had
the atomic weight of oxygen in Evanston air within 4-0.000,002 atomic
weight units.
B. Isotopic Composition of Photosynthetic Oxygen
As mentioned above, when the relatively rapid turnover of atmospheric,
water, and carbon dioxide oxygen is considered, one would have expected
atmospheric oxygen to have the isotopic composition of oxygen from the
photosynthesis process. A number of investigators (for a review see Kirshen-
baum, 34), including the author, have demonstrated that the oxygen liber-
ated in the photosynthesis reaction comes from the water rather than from
the carbon dioxide. Dole and Jenks (35) concluded that the liberated oxygen
had the isotopic composition to be expected if the following isotopic exchange
equilibrium between oxygen and water were established at the temperature
of and during the photosynthesis reaction:
O 1 6 t 2 tg) -k 2H~OlS(1) ~- 02 (g) -t- 2H2016(1) \ 18 (29)
The Soviet scientist, A. P. Vinogradov (36, 37), and his group of geo-
chemists believe, on the other hand, that the isotopic composition of the
oxygen liberated in the photosynthesis reaction has exactly the isotopic com-
position of water oxygen and that the very slightly enhanced O is content
of photosynthetic oxygen over that of the oxygen of water observed by Dole
and Jenks was due to concentration of O 18 in the gas phase due to plant
respiration during the experiment. Fractionation of oxygen isotopes on respira-
tion is discussed below. In Dole and Jenks's experiments the photosynthesiz-
ing algae were continuously illuminated while in the work of Vinogradov
et al. the illumination was natural with a dark period at night. During the
latter there would be relatively more respiration than in Dole and Jenks's
experiments. Whether the isotopic exchange equilibrium of reaction (29)
was established or not, the fact remains that the oxygen liberated in photo-
synthesis in ocean water has a smaller 018 content than atmospheric oxygen.
C. Isotopic Composition of Oxygen in Air Dissolved in Ocean Water
Rakestraw, R u d d , and Dole (38) made a careful study of the isotopic composi-
tion of the oxygen in air obtained by evacuation of ocean water taken at dif-
ferent depths in the ocean. Their data are illustrated in Fig. 2. A remark-
able correlation is seen between the decrease in the percentage of oxygen
in the air and the increase in the O 18 percentage. To quote Rakestraw, Rudd,
and Dole, " T h e conclusion is almost inescapable that marine vegetation,
2 0 S T R U C T U R E A N D F U N C T I O N O F O X Y G E N
calculated to be 1.009. There was considerable scatter to the data at low oxy- gen percentages due in part to the difficulty of making accurate O t s per- centage measurements in samples of air containing small concentrations of oxygen.
Vinogradov et al. (37) explain the above mentioned scatter in terms of
variations in temperature of the water at the four different geographical stations at which the dissolved ocean air was collected. Isotope fractionation factors usually change with the temperature. However, the average value of
a found by Vinogradov et al. (37) in their extensive work was 1.010 which
is within the limits of error equal to that found by Dole and coworkers, n a m e l y 1.009.
D. Theories of the Enhanced 018 Content of Atmospheric Oxygen
We have seen above that the oxygen produced in the photosynthesis reaction does not have the same isotopic composition as the oxygen already in the air, hence there must be some mechanism which tends to enhance the O 18 abun- dance in air. Calculations show that the atmospheric O 18/O16 ratio cannot be the result of the isotopic exchange equilibrium reaction because the equi- librium constant of the latter is m u c h too close to unity. Inasmuch as the oxygen of the air and the waters of the oceans covering most of the earth's surface are in continual contact, there cannot be any isotope exchange going on between the molecular species H~O and 02, otherwise the equilibrium would be established. In fact, there is no known mechanism by which H 2 0 and O5 can undergo oxygen isotope exchange in aqueous systems at room temperature; high temperatures and catalysts like platinum are required. Without taking space here to describe other possibilities which we have considered and rejected, let us turn at once to two reasonable possibilities. T h e fractionation of oxygen isotopes during the consumption of oxygen in the ocean suggested that this mechanism might explain the enhanced O is content of the atmosphere. Accordingly, Lane and Dole (39) investigated
tractionation during respiration of a n u m b e r of organisms, including Homo
sapiens. For the small organisms the apparatus used was that of Brown (40).
For Homo sapiens, oxygen was breathed in through a tube from a reservoir
of air whose oxygen supply was continually replenished, and breathed out through another tube. T h e exhaled breath passed through two large K O H bottles, then through a sampling flask before it was rebreathed. Two rubber balloons attached to the K O H reservoirs enabled the gas volume of the system to expand and contract with each breathing cycle and m a d e it pos- sible to make quick estimates of the a m o u n t of additional oxygen required. After a large fraction of the oxygen had been consumed, the isotopic com- position of the unreacted oxygen was determined in a mass spectrometer. T h e fractionation factor a was then calculated from the equation
MALCOLMDOLE, NaturalHistory of Oxygen 21 --m/ a O/^ __ y_^ - - y o e^ (31) Xo ~ x o e - m / a where y is the atom percentage of 0 is in the sample at the end of the experi- ment, y o is the 01 s percentage in the reservoir at the start of the experiment, Xo is the O is percengage in the oxygen added to the reservoir during the experiment, and rn is equal to the ratio of oxygen consumed to the a m o u n t initially present. If m u c h oxygen is respired so that m / a is equal to 3 or 4, the exponential terms can be neglected and a calculated directly from the ratio y / x o. T h e results of this work are illustrated in Fig. 3 where the dotted
I I '0--¢ I I I I [ © - ~ o I ooo I I I t i I I I I I I I I I I I I
~ 0 0 - - 0 0 ~ 0 [ ~ 0 ~ 0 ~ I I I LO0 1.02 a
HUMAN BEING
GREEN L E A V E ~
CRAB
FROG
VEGETABLES
FOREST L I T T E R
MUSHROOMS
MOLDS
BACTERIA
1.04 1.
FmtmE 3. Oxygen isotope fractiona- don factors observed during respiration or organic oxidation.
vertical line represents the fractionation factor necessary to account for the enhanced O is content of the atmosphere. As shown b y Dole, Hawkings, and Barker (41) the equation describing the changes per year in the oxygen isotope ratio of atmospheric oxygen is
N 1 - - No = An [ r ~ - ri] [l + Nd ?/o, where N1 = [018]/[016] in atmosphere at end of year. No = [018]/[01~] in atmosphere at start of year.
MALCOL~ DOLE Natural History of Oxygen 2 3
dioxide and molecular oxygen in an electric discharge, but that no exchange could occur between water vapor and oxygen. Vinogradov (37) has calculated that this mechanism could account for the enhanced 018 content of the atmosphere. T h e calculations given below, however, indicate that this proc- ess is too slow by about two orders of magnitude. F r o m radioactive fall-out measurements Vinogradov quotes 10 years as the time for the gases of the stratosphere to mix with those of the troposphere. Let us accept this estimate. T h e exchange between COs and Os can only be assumed to occur above the ozone layer because of the screening effect of the latter on solar ultraviolet radiation. If we adopt the 25 km height and above as the locus of the exchange, we have to realize that the COs-O~ oxygen a t o m "scrambling" can occur in only about 2 per cent of the atmosphere. T h e percentage of O ~8 in the stratosphere above 25 km as the result of all the oxygen in the COs becoming scrambled with the molecular oxygen of the atmosphere is readily calculated from the equation
per cent O is in stratosphere -- (0.2039) (0.998427) -t- (0.2078) (0.001573) (33)
where 0.998427 and 0.001573 are the total atom fractions in the mixture of oxygen in the Os and COs, respectively. In calculating the latter we have used the volume percentages of Os and COs in the atmosphere given by Glueckauf (43). T h e per cent 018 in the stratosphere after the exchange between all the CO2 and O~ is 0.203,906,134. In 1 year one-tenth of this oxygen will mix with all the oxygen below the 25 km level raising the per cent O ~8 in the troposphere as calculated by the equation
per cent 018 in troposphere after mixing with one-tenth of the COs-exchanged 02 = (0.2039) (0.998) q- (0.203,906,134) (0.002) = 0.203,900,012. (34)
Thus, the CO~-Os scrambling process should increase the 018 percentage in the troposphere by 0.006 ppm. In 1 year one-two thousandth or 0.0005 of the atmospheric O~. will be produced by the photosynthesis reaction. This will decrease the 018 per cent as calculated according to the equation
per cent 0 ~s in troposphere after I year's 02 production by photosynthesis = (0.2039) (0.9995) + (0.2003) (0.0005)
T h e difference between this n u m b e r and 0.2039 is 1.8 X 10 -6 or 0.9 ppm. Hence the photosynthesis process reduces the O 1s percentage by 0.9 p p m per year, but the COs-Os scrambling process raises it by only 0.006 p p m ; hence
2 4 STRUCTURE AND FUNCTION OF OXYGEN
the latter is an inadequate explanation of the enhanced O 18 content of the
atmosphere.
The theory of Lane and Dole (39) that fractionation of the oxygen isotopes
by respiration and organic oxidation causes the increase in the O 18 content
of the atmosphere seems to be the best from a quantitative standpoint.
E. The Oxygen Isotope Cycle in Nature
We can now set up the following steady state cycle of the oxygen isotopes in
nature:
Photosynthesis
/ /. ~ Greater O16 y i e l d - ~. ~ A
Land and tmosphere
Respiration and
organic oxidation
Greater O 1~consumption
In summary, photosynthesis yields oxygen containing a higher O16/O 18
ratio than the oxygen of the atmosphere, while respiration consumes oxygen
containing a higher O18/O I s ratio than the oxygen of the atmosphere and the
same ratio as that of photosynthetic oxygen, thus leading to the non-equi-
librium steady state value of the O16/O 18 ratio in the atmosphere. In other
words, the O le ratio of the atmosphere has been automatically adjusted to just
the right value so that the ratios for photosynthetic oxygen delivered to the
atmosphere and the oxygen extracted from the atmosphere by respiration
and organic oxidation are equal.
V I I. T H E F U T U R E O F A T M O S P H E R I C O X Y G E N
Berkner and Marshall (2) suggest that during periods of high rates of photo-
synthesis without m u c h organic matter to consume the oxygen produced,
the oxygen may have "overswung" its present level of abundance to a some-
what higher value. During the ice ages of the Permian period, the earth
cooled, photosynthesis was reduced "leading to a radical loss of atmospheric
oxygen." Thus, it is highly likely that the oxygen content of the atmosphere
has slowly fluctuated about a steady state concentration. Furthermore, in-
asmuch as the H atom escape must still be continuing, even if only at a very
slow rate, the net oxygen content of the atmosphere m a y be slowly rising.
Unfortunately, the problem of determining experimentally the atmospheric
oxygen percentage over millions of years in the past seems to be insuperable.
