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Studies on dissolved molecular oxygen in pure and sea water Mirhej, Michael Edward 1962

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STUDIES ON DISSOLVED MOLECULAR OXYGEN IN PURE AND SEA WATER by MICHAEL EDWARD MIRHEJ B.Sc, McGill University, 1958 M.Sc, University of Western Ontario, 1959 A thesis submitted in partial fulfilment of the requirements for the degree of DOCTOR OF PHILOSOPHY in the Department of CHEMISTRY We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA August, 1962 Tn presenting t h i s t h e s i s i n p a r t i a l f u l f i l m e n t of the requirements f o r an advanced degree at the U n i v e r s i t y of B r i t i s h Columbia, I agree tha t the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r reference and study. I f u r t h e r agree that permission f o r extensive copying of t h i s t h e s i s f o r s c h o l a r l y purposes may be granted by the Head of my Department or by h i s representatives. It i s understood that copying or p u b l i c a t i o n of t h i s t h e s i s f o r f i n a n c i a l gain s h a l l not be allowed without my written permission. The U n i v e r s i t y of B r i t i s h Columbia Vancouver 8, Canada. Department The University of British Columbia FACULTY OF GRADUATE STUDIES PROGRAMME OF THE FINAL ORAL EXAMINATION FOR THE DEGREE OF DOCTOR OF PHILOSOPHY of MICHAEL EDWARD MIRHEJ B.Sc, McGill University, 1958 M.Sc, University of Western Ontario, 1959 TUESDAY, OCTOBER 9, 1962, at 9:30 A.M. IN ROOM 342, CHEMISTRY BUILDING COMMITTEE IN CHARGE Chairman: F. H. SOWARD N. BARTLETT G. L. PICKARD M. BLOOM C. REID W. A. BRYCE P. M. WILLIAMS External Examiner: PROFESSOR PHILIP GEORGE Department of Chemistry University of Pennsylvania STUDIES ON DISSOLVED MOLECULAR OXYGEN IN PURE AND' SEA -WATER ABSTRACT Dissolved oxygen was studied in distilled water and salt solutions by means of nuclear magnetic resonance. The free induction technique was employed to measure the spin lattice relaxation time, T^ , for oxygen-free distilled water and for.water containing oxygen under one atmosphere of air and one atmosphere of oxygen at temperatures of 1° to 75° C. The same measurements were made for solutions of 0.5 M sodium chloride at temperatures of 1° to 40° C. The spin lattice relaxation probability, 1/T-^ C^\, in the presence of paramagnetic oxygen, was attributed to two relaxation mechanisms: the first a dipole-dipol interaction and the second a hyperfine interaction. The two terms were evaluated from measurements of Tl(c) a t t w o different magnetic field strengths at 20° C. Using the theoretical relationship between the dipole-dipole term and 1/T, the results were used to evaluate the hyperfine term at different temperatures. The activation energy obtained from the variation of the dipole-dipole term with temperature was assumed to represent the energy required to break one hydrogen bond between two water molecules. The activation energy found for the hyperfine term was taken as a measure of the breaking of a hydrogen bond between one oxygen molecule and an aggregate of water molecules. The activation energy of the hyperfine term was found to be a function of temperature. Comparison of this quantity with heats of solution of dissolved oxygen in distilled water and salt solution showed a similar pattern of change in both. Oxygen supersaturation was studied in sea water cultures of Nitzschia closterium and Chlorella strain "A" at a temperature of 12° C. Saturation values up to 2007« were reached under illumination with light energy of 9.2 x 10"langlies/min. Nitzschia was found active, under the same Chlorella. to be more photosynthetically culture conditions, than Oxygen production by Nitzschia was shown to be a function of the difference in photosynthetic pigment concentrations (chlorophyll-a — — non-astacin carotenoid), total alkalinity, and the change in catalytic activity of the medium. Variation of oxygen concentration in Nitzschia cultures under light and dark periods indicated a mechanism by which oxygen may escape as microbubbles to the atmosphere. The rate of oxygen desupersaturation was measured in water free of organisms. The rate increased with increase of ion content and with the surface to volume ratio of the water column, but was not influenced by addition of siliceous particulate matter to the supersaturated water. Small addition of a surface active agent (heptanoic acid) increased the oxygen desupersaturation rate but further addition decreased the rate. GRADUATE STUDIES Chemistry Seminar in Physical Chemistry Advanced Chemical Oceanography Topics in Inorganic Chemistry Surface Chemistry Physical Organic Chemistry Synthetic Methods in Organic Chemistry D.E. McGreer P.M. Williams H.C. Clark W.R. Cullen N. Bartlett J. Halpern R. Stewart R.A. Bonnett D.E. McGreer W. McRae Oceanography Introductory Courses in Oceanography Oceanographic Methods Phytoplankton and Photosynthesis G.L. Pickard R.F. Scagel B. McK.Bary G.L. Pickard R.F. Scagel i i ABSTRACT Dissolved oxygen was studied in d i s t i l l e d water and salt solutions by means of nuclear magnetic resonance. The free induction technique was employed to measure the spin lattice relaxation time, T-^ , for oxygen-free d i s t i l l e d water and for water containing oxygen under one atmosphere of air and one atmosphere of oxygen at temperatures of 1° to 75°C. The same measurements were made for solutions of 0.5 M sodium chloride at temperatures of 1° to 40°C. The spin lattice relaxation probability, 1/T^^, in the presence of paramagnetic oxygen, was attributed to two relaxation mechanisms: the f i r s t a dipole-dipole interaction and the second a hyperfine interaction. The two terms were evaluated from measurements of T^ c^ a t t w o different magnetic f i e l d strengths at 20°C. Using the theoretical relationship between the dipole-dipole term and ^/T, the results were used to evaluate the hyperfine term at different temperatures. The activation energy obtained from the variation of the dipole-dipole term with temperature was assumed to represent the energy required to break one hydrogen bond between two water molecules. The activation energy found for the hyperfine term was taken as a measure of the breaking of a hydrogen bond between one oxygen molecule and an aggregate of water molecules. i i i The activation energy of the hyperfine term was found to be a function of temperature. Comparison of this quantity with heats of solution of dissolved oxygen in d i s t i l l e d water and salt solution showed a similar pattern of change in both. Oxygen supersaturation was studied in sea water cultures of Nitzschia closterium and Chlorella strain "A" at a temperature of 12°C. Saturation values up to 200% were reached under illumination with light energy of 9.2x10 langlies/min. Nitzschia was found to be more photosynthetically active, under the same culture conditions, than Chlorella. Oxygen production by Nitzschia was shown to be a function of the difference in photosynthetic pigment concentrations (chlorophyll-a — non-astacin carotenoid), total alkalinity, and the change in catalytic activity of the medium. Variation of oxygen concentration in Nitzschia cultures under light and dark periods indicated a mechanism by which oxygen may escape as microbubbles to the atmosphere. The rate of oxygen desupersaturation was measured in water free of organisms. The rate increased with increase of ion content and with the surface to volume ratio of the water column, but was not influenced by addition of siliceous particu-late matter to the supersaturated water. Small addition of a surface active agent (heptanoic acid) increased the oxygen desuper saturation rate but further addition decreased the rate. iv TABLE OF CONTENTS Page PART I - PROTON SPIN RELAXATION BY PARAMAGNETIC MOLECULAR OXYGEN 1 Introduction 1 Theory „ 14 Experimental 19 I. Apparatus 19 II. Reagents , 20 III. Procedure 21 Results and Discussion . . «, 26 I. (i) Determination of T^, 26 ( i i ) Behaviour of with oxygen concentration 30 II. Relaxation terms 35 (i) Evaluation of the relaxation terms 41 ( i i ) Comparison of activation energies and heats of solution 43 General Discussion 52 Conclusion 58 Suggestion for Further Work 58 PART II - ASPECTS OF OXYGEN SUPERSATURATION IN SEA WATER. . . 60 Introduction 60 Experimental 76 I. Methods and apparatus 76 V Page II. Reagents 80 III. Procedure 81 Results and Discussion 86 A (i) Variation of oxygen and pH 86 A ( i i ) Relation of growth and photosynthetic activity to oxygen supersaturation . . . . 89 a. Number of organisms 89 b. Photosynthetic pigment concentration . . . 95 c. Total alkalinity 99 d. Bacterial contamination 100 A ( i i i ) Relation of catalytic activity to oxygen production 102 B Effect of light and dark periods on oxygen concentration in a heavy culture 106 C (i) Influence of water column geometry on rate of desupersaturation 110 C ( i i ) Influence of salinity on rate of escape of oxygen 113 C ( i i i ) Influence of diatomaceous earth and surface tension on desupersaturation rate constant 114 Conclusion 118 References 156 LIST OF TABLES Table Page I. T x (Measured) 27 II. Values of T^ at Different Temperatures and Concen-trations of Oxygen - D i s t i l l e d Water 28 III. Values of T x at Different Temperatures and Concen-trations of Oxygen - 0.5M Sodium Chloride 29 IV. Values of 1/T-^  at Different Oxygen Concentrations -D i s t i l l e d Water 31 V. Values of l/T^ at Different Oxygen Concentrations -0.5M Sodium Chloride 31 VI. 1/T-L Values of Oxygen-Free and 32.00 mg./l. Oxygen-Concentrated D i s t i l l e d Water 33 VII. 1/TX Values of Oxygen-Free and 32.00 mg./l. Oxygen-Concentrated 0.5M Sodium Chloride 34 VIII. Evaluation of C r f c Values at Different Temperatures - D i s t i l l e d Water 42 IX. Evaluation of C T c Values at Different Temperatures - 0.5M Sodium Chloride 42 B B X. Values of ^  g rr" a n d 7p~ a t Different Temperatures -D i s t i l l e d Water 44 •p -p XI. Values of -—7 and -~- at Different Temperatures -W S e ^e 0.5M Sodium Chloride 45 v i i Table Page XII. Estimation of the Activation Energy of T c -D i s t i l l e d Water 47 XIII. Estimation of the Activation Energy of ^ -0.5M Sodium Chloride 48 XIV. Ratio of the Relaxation Terms and Estimation of the Activation Energy of Tg - D i s t i l l e d Water 49 XV. Ratio of the Relaxation Terms and Estimation of the Activation Energy of T £ - 0.5M Sodium Chloride . . 49 XVI. Evaluation of Heats of Formation for Solubility of Oxygen under One Atmosphere of O2 - D i s t i l l e d Water 50 XVII. Evaluation of Heats of Formation for Solubility of Oxygen under One Atmosphere of O2 - 0.5M Sodium Chloride 50 XVIII. Heat of Solution and Activation Energy of Dissolved Oxygen in Water 51 XIX. Variation of O2, N2 and pH in Cultures of Nitzschia and Chlorella 87 XX. Series Culture of Nitzschia 90 XXI. Quantities Influencing Oxygen Production in Nitzschia Culture 92 XXII. Variation of Oxygen Concentration with Catalytic Activity - Nitzschia 104 v i i i Table Page XXIII. Variation of Oxygen Concentration with Catalytic Activity - Chlorella 104 XXIV. Variation of Oxygen Concentration under Light and Dark Periods in Nitzschia Culture 107 XXV. Effect of Surface to Volume Ratio on Desupersatura-tion Rate Constant. S = 30.2% 112 XXVI. Effect of Surface to Volume Ratio on Desupersatura-tion Rate Constant - D i s t i l l e d Water 112 XXVII. Variation of Desupersaturation Rate Constant with Salinity under Stirred Conditions 114 XXVIII. Influence of Diatomaceous Earth Wall Effect on Kdes 115 XXIX. Influence of Surface Tension on Kdes 115 ix LIST OF FIGURES Figure Page I. Proposed Dipole Orientation of a Water Molecule in Presence of Ions 120 II. (a) NMR Assembly 120 II. (b) Induction Coil Chamber 120 II. (c) Degassing Instrument 121 II. (d) Sample Tube 121 III. Oscilloscope Display of the Two Pulses and the Detected r - f Signal 121 IV. Determination of T^ 122 V. Smoothed Curves of T^ Values at Different Temperatures (Disti l l e d Water) 123 VI. Smoothed Curves of T^ Values at Different Temperatures (Salt Solution) 124 VII. Comparison of T^ Values for Oxygen-Free D i s t i l l e d Water 125 VIII. Variation of 1/T^ Values with Oxygen Concentrations (Dis t i l l e d Water) 126 IX. Variation of 1/T^ Values with Oxygen Concentrations (Salt Solution) 127 X. Behaviour of 1/Ti( c) with Temperature (Di s t i l l e d Water) 128 XI. Behaviour of 1/T^(C) with Temperature (Salt Solution) 129 X Figure Page XII. Change of C T c Term with Temperature 130 •p XIII. Change of — Y r r ~ T e r m w i t n Temperature 131 e XIV. Determination of the Activation Energy E c 132 XV. Determination of the Activation Energy E e at Different Temperatures (Distilled Water) . . . . .133 XVI. Determination of the Activation Energy E e at Different Temperatures (Salt Solution) 134 XVII. Determination of the Heat of Solution - &H for Oxygen in Water under 1 Atm. Pressure 135 XVIII. Variation of the Ratio of the Two Relaxation Terms with Temperature 136 XIX. Two Different Orientations of Two Dipole Molecules .137 XX. Variation of Oxygen and pH in Nitzschia Culture with Time 138 XXI. Variation of Oxygen and pH in Chlorella Culture with Time 139 XXII. Comparison of Number of Organisms in a Nitzschia Culture with Photosynthetic Pigment Concentration 140 XXIII. Comparison of Number of Organisms in a Nitzschia Culture with Rates of Photosynthesis and Respiration 141 XXIV. Behaviour of Photosynthetic Rate in Nitzschia with the Difference of Pigment Content 142 x i Figure Page XXV. Behaviour of Net:Gross Photosynthetic Rates in Nitzschia with Ratio of Pigment Content 143 XXVI. Dependence of Respiration Rate in Nitzschia on Carotenoid Pigment Concentration 144 XXVII. Relation of Per Cent Saturation to Rates of Photosynthesis and Respiration in Nitzschia. . . .145 XXVIII. Influence of Total Alkalinity on the Saturation Level in a Nitzschia Culture 146 XXIX. Variation of Oxygen and Catalytic Activity in a Nitzschia Culture with Time 147 XXX. Variation of Oxygen and Catalytic Activity in a Chlorella Culture with Time 148 XXXI. Extracellular Carbohydrate Relative to Number of Organisms 149 XXXII. Change of Oxygen Concentration under Illumination in a Nitzschia Culture at Different Depths . . . .150 XXXIII. Change of Oxygen Concentration under Dark in a Nitzschia Culture at Different Depths 151 XXXIV. Variation of Oxygen Desupersaturation Rate Constant with Surface to Volume Ratio of a Water Column . .152 XXXV. Variation of Oxygen Desupersaturation Rate Constant with Salinity under Stirred Conditions 153 x i i Figure Page XXXVI. Behaviour of Siliceous Surface Area with Oxygen Desupersaturation Rate Constant 154 XXXVII. Influence of Heptanoic Acid Addition on Oxygen Desupersaturation Rate Constant 155 x i i i ACKNOWLEDGEMENTS The author wishes to express his thanks and gratitude to Dr. M. Bloom for his assistance and guidance as well as his provision of experimental equipment . Thanks are also extended to Doctors G. L. Pickard, C. Reid, R. F. Scagel and P. M. Williams for their suggestions and criticism. The measurement of diatomaceous earth surface, courtesy of Dr. I. H. Warren, and the supply of phytoplankton strains by Dr. R. Guillard of Woods Hole Oceanographic Institution and Dr. B. Sweeney of Scripps Institution of Oceanography are greatly appreciated. The financial assistance of the National Research Council is gratefully acknowledged. PART I PROTON SPIN RELAXATION BY PARAMAGNETIC MOLECULAR OXYGEN INTRODUCTION The aim in this part is to study the effects of dissolved oxygen on the spin relaxation of protons in d i s t i l l e d water and in salt solutions, at different concentrations of oxygen and different temperatures, by means of Nuclear Magnetic Resonance. It is also the aim to infer from, and correlate with the results any factors that bear on the solubility of oxygen in both dis-t i l l e d and ion-containing water, and to show in particular the close association between the degree of solubility and the tendency of the water solvent to form aggregates and clusters. The solubility of a gas in a liquid such as oxygen in water, where Henry's law applies (Henry, 1803), is governed by the relation that the partial pressure of the gas in equili-brium with i t s solution in a liquid i s found to be proportional to i t s concentration in solution: P 2 = KC 2 (1) The proportionality constant, K, depends on the temperature, and on the nature of the solute and solvent, and is independent of pressure. Notable treatments for this kind of behaviour have been 2 carried out thermodynamically and kinetically (Moelwyn-Hughes, 1957) on the basis of experimental results obtained from solubilities and their variation with temperature. The thermo-dynamic approach was used to determine the chemical potential of the solute in a solution of unit concentration, and then further employed to explore the free energies of solutions from which many of their properties are derived. The kinetic treatment, on the other hand, was based on a certain rate of condensation of solute from the gas phase, and on i t s evaporation from the liquid, a l l taking place at an uppermost, thin layer in the liquid, and on the probability that the solute and solvent molecules possess a certain energy between them, which i f exceeded could eject the solute molecule from that particular surrounding. The nature of bonding, and the physical picture of the micro-structure of solvent and solute was not dealt with u n t i l recently. Heidt and Johnson (1957) advanced a theory of molecular oxygen hydrates in water based on studies of ultra-violet absorption at different temperatures. They postulated the existence of at least two hydrate species existing at equilibrium with each other, and attributed such occurrence to hydrogen bonding. The model suggested was a new kind of a hydrogen bridge which takes place between the diradical form of molecular oxygen and water or paraffin hydrocarbons. In this 3 bridge the proton would be shared by three electrons, one of which comes from the diradical form of the oxygen molecule and the other two from the original electron pair bond 0-H in the case of water and C-H in the case of paraffin hydro-carbons. In this kind of bridge the proton would be nearer to the 0 or C atom, respectively, of i t s original electron pair bond than to the oxygen atom of molecular oxygen. The resulting structure would be a ring in the case of 1:1 complex O^ '.H^ O and a chain in the 1:2 complex 0^:2^0. Although this appears to be a feasible approach, nonetheless, the treatment gives an oversimplified picture which excludes effects of other processes taking place in liquid water. These processes include complex formation and the tendency of water to form aggregates, which may have a great influence in determining the extent of oxygen solubility in water. The present concepts of water aggregates are based mainly on information obtained from x-ray, neutron diffraction, Raman spectra and infra-red studies of ice (Lonsdale, 1958; Ockman and Sutherland, 1958; Frohnsdorff and Kingston, 1958) and water (vanEck et a l . , 1958; Frank, 1958; Brockhouse, 1958). Hydrogen bonding between water molecules is found to play a leading role in a l l the interpretations advanced for the struc-ture of liquid water and aqueous solutions. This stems in effect from the structure of a free water molecule which has 4 it s two hydrogens at 0. from the oxygen atom, with the angle between the 0-H bonds ~ 105°. In ice, 0-H bonds on one molecule point towards the oxygen of adjacent molecules, forming a structure in which four hydrogens are arranged nearly tetra-hedrally around each oxygen, two being at 1. oiX and two at 1.75% (Peterson and Levy, 1957). Thus, water is pictured as a broken-down, ice-like struc-ture which has the capacity to form clusters that are hydrogen bonded to each other. The bonds break and form continuously with the motion of the molecules in the liquid. The continuity of hydrogen bonding, however, i s in question. Lennard-Jones and Pople (LJP) (1951), use the term hydrogen bond to refer to any of a wide range of 0-H 0 configurations, calling the bond "bent" when either the lone pair or the H atom has moved out of the 0-0 line as a result of rotation of i t s water molecule. They picture rather extreme bending as being possible without the need to give up the term hydrogen bond. In contrast to this, Frank and Wen (1957) distinguish qualitatively between large and small angles of bend, assigning the former to non-bonded and the latter to bonded categories. A "small" angle of bend means an angle small enough so that the angular require-ments for tetrahedral covalent bonding are not transgressed. From this latter hypothesis liquid water emerges as con-sisting of flickering clusters of bonded molecules mixed with, 5 and alternating roles with, non-bonded f l u i d which encloses them and constitutes the rest of the sample. A (larger or smaller) cluster is pictured as forming when the stage i s set by an energy fluctuation which creates a suitably "cold" region, and relaxing when the necessary energy of "melting" becomes available. This flickering cluster representation receives the support of Wang e_t a l . (1953) since i t offers an explanation of the energy of activation in liquid water being 4.6 kcal/mole whether calculated for dielectric relaxation, self-diffusion or viscous flow. This is just what would be expected i f the principal requirement for any of these processes in water were the breaking down of a rather r i g i d structure. The work of Eigen and DeMaeyer (1958) on the protonic charge transport in water presents a similar picture to that of cluster formation. In this case three water molecules 4- + surround a H^ O ion and form a H^ O^  complex. The charge transfer inside the complex is very fast and of the order of - 1 3 -14 10 to 10 s e c , while the charge transfer to a molecule outside the complex is governed by the formation of a new hydrogen bond. The average time of liberation of water mole-cules from the complex and H-bond formation at the complex can be related to the dielectric relaxation times which also involve a mechanism of liberation from the structure and orientation of the dipole-dipole. The slowest dielectric relaxation process in water is characterized by a time constant of about 10~^ sec. and an activation energy of 4.5 kcal/mole (Lane and Saxton, 1952). This process corresponds to the liberation of double H-bonded water molecules. Hindman (1961) in his investigation of NMR effects in aqueous solutions assigns a value for the number of hydrogen bonds being ruptured in liquid water at some temperature from data based on those of Haggis et a l . (1952). He estimates 1.82 hydrogen bonds per water molecule at 0°C, corresponding to 9% of the bonds being broken in going from ice to water. This value is between Pauling's (1960) estimate of 15% of the bonds being broken at 0°C. and Pople's (1951) estimate that very few bonds are broken. In the light of such interpretations, the addition of ionic and non-polar species to water should be examined, and their effects on the clusters or aggregates of water be taken into consideration i f any explanation for the difference in oxygen solubility between pure and salt water solutions at different temperatures is to be attempted. The addition of ionic salts, such as NaCl, has been known (Frank and Wen, 1957) to result in a structure-breaking effect of the clusters on one hand, and on the other hand to form water of hydration which is irrotationally bonded with a specific hydration number for each certain ion. Frank and 7 Evans (1945) suggest that essentially three regions exist around a dissolved ion: (1) a "frozen" layer of water oriented by the ion, i.e., the hydration shell; (2) an intermediate disorganized region caused by the i n a b i l i t y of the hydrated ion to f i t into normal water structure; and (3) the region of normal water structure. For the ion-water complex, or the hydrated shell, two principal models exist. One i s based on detailed energy considerations and was advanced by Bernal and Fowler (1933) and the other by Verwey (1942). The Bernal-Fowler model assumes a cation would bind a water molecule in coplanar fashion with the negative end of the water dipole directed toward the ion as shown in Figure I (a), whereas an anion would be preferentially oriented with one hydrogen pointing toward the ions as in Figure I (c) rather than the dipolar position in (d). A water bonded as in Figure I (a) could bond to only two additional water molecules and hence one hydrogen bond must be broken per coordinated molecule, whereas a water bonded as in Figure I (c) could bond to three additional water molecules and hence one bond must be broken for each two coordinated water molecules. Verwey, on the other hand, suggests that univalent cations are attached to the region of negative charge due to the lone-pair electrons of the oxygen in an angular tetrahedral fashion Figure I (b), such that 8 the coordinated water molecule i s able to bond to three additional water molecules. The d i e l e c t r i c constant data, however, gives evidence for difference i n the modes of bonding of water around anions and cations (Hindman, 1961). The lowering of the d i e l e c t r i c constant has been interpreted i n terms of a model i n which water oriented as i n Figure I (a) i s not free to rotate, while the n e g l i g i b l e e f f e c t of anion on the d i e l e c t r i c constant has been interpreted as due to free r o t a t i o n of water molecules bonded as i n Figure I ( c ) . Thus the Bernal-Fowler model i s considered to be more reasonable than the Verwey one. In resume then, addition of ions to water 1. disrupts the water structure and breaks hydrogen bonds; 2. the i o n i c e l e c t r i c f i e l d d i s t o r t s the e l e c t r o n i c structure of water molecules; 3. the water of hydration i s dependent on the size of the cation; and 4. three d i s t i n c t s i t e s for water molecules are provided: near cation, near anion, near other water molecules (Fabricand and Goldberg, 1961). The addition of non-polar solutes tends to replace the disordered neighbours around the edges of the c l u s t e r . The motion of the disordered neighbours provides the mechanism by which the clu s t e r receives the "heat of melting" which 9 ultimately terminates i t s existence. If some of these neighbours are replaced by non-polar solute molecules or non-polar groups attached to solute molecules, the cluster boundary w i l l be stabilized, for i t s new neighbours are only weakly able to transmit displacements and torques (Frank, 1958). Also, i t is expected for a solute such as an oxygen molecule to be always found at the extremities of the cluster i f bonded to one or more water molecules. The properties of an oxygen molecule hinder i t s f i t t i n g in the middle of the cluster because of i t s ina b i l i t y to arrange i t s e l f in a tetrahedral form in the same manner as water molecules. This also presents the notion that any complex formation, or breaking, taking place between water and oxygen must occur in a jump process where the oxygen molecule w i l l break from the edge of one cluster and hop to the edge of the other when they are approximately at a lat t i c e distance away, or on the "melting" of one cluster the oxygen molecule finds i t s e l f sticking to another being formed. Any postulate of complex formation between oxygen and water molecules should take into account the structure of the oxygen molecule and i t s capability to participate in hydrogen bond formation. The molecular orbital expression for two oxygen atoms forming an oxygen molecule gives the following electronic arrangement (Cartmell and Fowles, 1961): 0 2 KK(z(T) 2(y<r) 2(x<r ) 2(w TT) 4(vIT) 2 Here KK designate the inner (ls) electrons, (x<r) and (WIT) o the sigma and p i bonds formed between the two atoms, (z<r) 2 and (ycr) the two unshared pairs of electrons at both ends 2 of the molecule along the bond axis, and (VTT) the two anti-bonding electrons which give the molecule i t s paramagnetic properties. Since any postulate of hydrogen bonding with any specific orbital w i l l have to take into consideration the hybridization, and derealization of the orbital, such treatment w i l l not be attempted and w i l l be considered beyond the scope of this work. However, the oxygen molecule as a whole w i l l be considered in terms of polarizability from an electrostatic point of view. This qualitative approach would account for heats of solution in terms of the electrostatic interaction between the oxygen molecule and the dipole moment of water, taking into account the numbers of bonds broken or formed at different temperatures, and the effect of ions on both solvent and solute. Further electrostatic approach w i l l be examined as a function of the energy of point charge dipoles, and of the distance between the points as well as the dependence on the angle of approach (Pimentel and McClellan, 1960). The paramagnetic property of the oxygen molecule, and i t s capability to dissolve in pure and ion-containing water, makes the solution an interesting subject for study by means 11 of nuclear magnetic resonance. Previous studies of proton relaxation by paramagnetic ions in aqueous solutions (Solomon, 1955; Bloembergen, and Morgan, 1961) provide a lead and a theory for such investigations. Although the solubility of molecular oxygen at ordinary temperature and atmospheric 17 -3 pressure is of the order of 2x10 molecules cm , the effect on T^, the proton spin-lattice relaxation time, is appreciable and sufficient to produce observable changes from those of oxygen-free water. Guilloto's (1956) work clearly showed changes of T^, the spin-lattice relaxation time, at 20°C. , with different concentrations of oxygen in solution, and the linear dependence of 1/T^ on concentration. This was similar in manner to the linear relaxation effect of paramagnetic -H- -H- 15 concentrations of Ni and MM ions in a range from 2x10 to 2 x l 0 1 8 ions cm"3 (Hennel et a l . , 1961). Earlier studies of paramagnetic impurities in water were interpreted assuming that the relaxation was due only to dipole-dipole interaction between electron and nuclear spins, the fluctuations of these interactions being governed by viscosity and diffusion. Bloembergen et a l . (1948) found the addition of paramagnetic ions to water decreased both the spin-la t t i c e and the spin-spin relaxation times, T^ and The additional contribution to the relaxation rate was found to be proportional to the concentration. 12 Later, Conger (1953) and Morgan et a l . (1956) concluded from measurements of relaxation times in water-glycerol. solu-tions of several aquated ions that the observed changes were not entirely attributable to the change in viscosity according to the early theory. It was suggested that the difference observed might well provide an insight into the nature of molecular interactions in those and similar solutions. Morgan suggested that at least part of the difference lay in the presence of the non-exchangeable protons of the glycerol mole-cules . Bloembergen (1957) followed with comparison of the ratio of T^/T^ in paramagnetic ion solutions finding a deviation -H-from unity for most ions to as high as 7.1 for Mn . The interpretation advanced emphasized that a spin exchange inter-action was taking part in the relaxation mechanism, and this was due to the formation of a hydrated complex between the paramagnetic ions and water molecules. This was supported and extended by Bloembergen and Morgan (1961) to interpret the different effect on relaxation between various paramagnetic ions and the role played by the spins of the ions, the geometry and number of their water of hydrations as well as their relative motions in solution. A similar approach to the above was employed here for investigating the nature of the solubility of molecular oxygen to provide information on complex formation and hydrogen bonding. 13 This should be especially informative i f the complex included an oxygen molecule and a water aggregate in cluster, since under this condition a spin exchange interaction would take place i f the amplitude of the wave function of the unpaired C«2 electrons at the protons was non-zero. The effect on the dipole-dipole interaction on the other hand i s related to the tumbling of the water cluster and this in turn is governed by temperature and viscosity. The addition of one oxygen molecule to the cluster w i l l be assumed not to alter the dynamical motions of the complex since the size of the cluster i t s e l f is considered to be much larger than the size of an oxygen molecule. Changes in viscosity may play some role in nuclear spin relaxation and at the same time may be correlated with formation or breaking of water clusters. The addition of non-paramagnetic ions such as Na + and Cl to water could provide significant relaxation time effects, via a mechanism in which the ions would obtain their hydration spheres and break the water structure without contributing any paramagnetic influence. Since the solubility of oxygen is also decreased by the addi-tion of such ions, the contrasts and similarities between influences of solubilities and influences on relaxation times due to changes in complex formation should become apparent. 14 THEORY PROTON SPIN RELAXATION BY PARAMAGNETIC 0 2 The proton spin relaxation in presence of molecular oxygen is assumed to be due to two kinds of interaction. The f i r s t i s a dipole-dipole interaction between the proton and electron spins, and the second a proton-electron spin exchange or hyperfine interaction due to the formation of a complex between water and oxygen molecules. The Hamiltonian for the interaction between the proton on ^ 0 and the paramagnetic electrons of 0 2 could then be described as H » H dip-dip + A I . S (2) where A i s a scalar constant; I the nuclear spin system; and -» S the electron spin system. The interaction (2) produces transitions between the nuclear Zeeman levels and the electronic spin Zeeman level and exchanges the Zeeman energy with the energy of the lat t i c e or both. The (21+1) neighbouring Zeeman levels for a nuclear spin I are separated by t\ CO^ . = ^j1r\ H q in an external f i e l d Hq, where ^1 is the nuclear gyromagnetic ratio, while for an electron spin S, the electron spin Zeeman levels are s p l i t by ^ CJg = ^s^^o' ^S being t n e electron gyromagnetic ratio. For protons = 2.67x10^ while for electron spins on the 0 2 molecules V e = 1.75x10^, so that CO„ O 15 When the nuclear spin system is in equilibrium with the l a t t i c e , the distribution of spins in the Zeeman levels is described by a Boltzman distribution for temperature T, the temperature of the l a t t i c e . The interest here is with the approach of the nuclear spin temperature towards equilibrium after i t has been disturbed by a radio frequency pulse. This approach is governed by a time constant T^, the spin-l a t t i c e relaxation time. The contribution of the interaction with the S system given by equation (2) to the proton spin relaxation probability l/T., , . is given by Abragam (1961; equations 88 and 120, 1(c) Chap. VIII). By evaluating the correlation functions of both terms (Bloom, 1962), i t may be noted that the dipole-dipole term and hyperfine term have different correlation times, T and T respectively, associated with them. This is c e so because the dipole-dipole interaction depends on the vector -s -+ -» y joining the proton-O^ pair, while the scalar I . S inter-action is independent of the orientation of this pair. If each 0 2 molecule is assumed to be interacting with a complex containing n^ protons which behaves for some time as a r i g i d unit and i f the correlation functions are assumed to have an exponential time dependence, the general expression for T, , v is as follows: 1(c) n ^ „ 2 , . ( /.) cr a. Z Tc ] TT" < E dd> \ (_ + — r — - r n u 0) s "J S c 3 i+<y2cf 2 \ h 9 +i 9 v ^ 9 ^To "I + f S(S+l ) ~ r < A > \ *i-Zo(n. 3 /c°s ' i + o 2 T 2 \ h S e (3) Where <E 2dd> = % l S(S+1) < > (4) is the I S r" average value of the dipole-dipole interaction with the protons 2 of the complex; A is the average value of the hyperfine interaction in the complex; n the number of paramagnetic 3 1 1 1 molecules per cm of water; and T l ( c ) T l ( 0 2 ) Tl(pure) where T., , . i s the spin-lattice relaxation time in oxygen l(pure) free water and T^^ 02) *"S t* i e sP*-n relaxation time in presence of oxygen in water. Three additional approximations have been made in writing equation (3). (a) C0 g » as stated earlier so that ^ 2 2 COc - CO (b) « !• No restriction i s made on u) T o r ^ J and, (c) The spin temperature of the o c D e S system of the electron is taken to be the same as the l a t t i c e temperature. That is in the absence of any perturbing r - f at the Larmor frequency £0g of the S system, this condition is f u l f i l l e d i f T^g^ T 1 ( I ) ' S U c h c o n d i t i o n i s normally satisfied since y g ^ so that the electronic spins are very strongly coupled to the lat t i c e by comparison with the nuclear spins. Equation (3) i s i d e n t i c a l with the expressions developed for relaxation of protons i n aqueous solution of paramagnetic ions (Solomon, 1955; Bloembergen and Morgan, 1961), and which proved to be the most successful form for c o r r e l a t i o n of experimental r e s u l t s . The model envisaged i n the present work i s of an oxygen molecule coming i n touch with a water clu s t e r for a time H"^. The complex formed then rotates and tumbles as a whole i n the body of the l i q u i d with a c h a r a c t e r i s t i c time T • I t should also be noted that i f the 0^ spins are relaxed i n a time T g , the electron spin relaxation time, then the values of T c and T are as follows: e and s i m i l a r l y , In general T c w i l l be a function of temperature and magnetic f i e l d , while T r and w i l l be functions of tempera ture only. The r e l a t i v e magnitudes of the three c o r r e l a t i o n times I T , , fT and r r have been investigated for paramagnetic ions h r o i n aqueous solutions where i t was found that and are much larger than T . Bloembergen and Morgan ( 1 9 6 1 ) have suggested that for such a case T c " T c ° exp (E c/RT) (7). In this work i t is also assumed that 0* is mainly depen-c dent on T , and that col T 2 « 1, and CJ 2 T 2 » 1. That r D C b e is T is much less than 10 sec. and T is much larger c e —11 than 10 sec. at the value of (0 at 7000 gauss, the magnetic D 11 ""I f i e l d used in this experiment where ^  = 1.24x10 sec. Equation (3) could then be resolved into the following form: = C T + — ~ (8) T K c ) c 6)? rr s where C - | ^ n <E2dd> i | n h ; and B - | S(S+l)nn h < A2> . 19 EXPERIMENTAL I. Apparatus (i) Nuclear Magnetic Resonance by Free Induction Technique: The apparatus used was a conventional pulsed nuclear magnetic resonance apparatus (Hahn, 1950) and consisted of a magnet with a power source connected to a current stabilizer and operated at 7000 gauss f i e l d strength; a pulsed oscillator adjusted to the Larmor precession frequency of 29.80 Mc for the protons; a pulse timer; a trigger; 30 Mc amplifier and detector arranged as in Figure II (a). The signal height was read on an oscilloscope at an output of 2 volts/cm. ( i i ) Thermostated Induction Coil Chamber: A copper c o i l %" diameter was wound in a cylindrical form to a height of 6 cm. and a diameter of 3 cm. This was insulated with aluminum f o i l on both sides and covered with Scotch e l e c t r i c a l tape on the outside, leaving a small opening at the top for the sample entrance. The c o i l was placed ver t i c a l l y on a wooden board around the induction c o i l as in Figure II (b). The ends of the copper c o i l were connected to a circulating water bath with a temperature range from 1°C. to 90°C. and temperature control of * 0.1°C. A thermocouple was placed inside the chamber to determine temperature fluctua-tions at a set temperature. Fluctuations were of the order of T O T O - 0.5 C. at low temperatures and - 1.0 C. at high temperatures. ( i i i ) Degassing Instrument: A one l i t e r bulb with openings at both ends was sealed to a vacuum system on one side, and to the sample tube on the other. The bulb contained a teflon-covered bar magnet 2 cms. in length, and was placed over a magnetic st i r r e r as in Figure II (c). (iv) Sample Tubes: These were of pyrex glass 0.6 cm. I.D. and about 8 cm. in length. Two designs of sample tubes were used, one for the sample of oxygen-free water which was sealed and used permanently, and the other for oxygenated water under either one atmosphere of air or one atmosphere of oxygen (Figure II (d) ). II. Reagents (i) Double d i s t i l l e d water was used in a l l samples: oxygen-free, oxygen-concentrated, and salt solutions. ( i i ) Compressed Air: "Liquid Air" cylinder with 21% 0^ concentration was used; in addition to 0^, the gases in mixture were the same as atmospheric gas content. ( i i i ) Compressed Oxygen: "Liquid Air" cylinder containing 99.5% oxygen was used; the rest of the gas was nitrogen (0.5%). 21 (iv) Sodium Chloride Crystals: "B & A" reagent grade, which contained metal impurities of 0.00057o as lead, and 0.0002% iron. IEU Procedure Spin l a t t i c e relaxation times were measured for oxygen-free double d i s t i l l e d water, and for water containing oxygen under one atmosphere pressure of a i r , and one atmosphere pressure of oxygen, at temperatures of 1° to 75°C. The same measurements were carried for solutions of 0.5M sodium chloride at tempera-tures from 1° to 40°C. (i) Degassing Method: (a) Double d i s t i l l e d water: 150 ml. of water was placed in the degassing bulb, Figure II (c). Both stopcocks A and B were opened and the mercury diffusion pump, o i l pump and magnetic stirrer were switched on. Degassing was carried out for one-half hour, after which stopcock A was closed, then B. The water in the bulb was t i l t e d u n t i l three or four ml. gathered in section C; the bulb was then warmed by hand to increase the vapor pressure. Stopcock A was opened and the water was displaced under pressure to the sample tube. The sample was sealed at the constriction and used thereafter in relaxation time measurements, (b) 0.5M Salt Solution: 28.0 g. of sodium chloride was dissolved in 1000 ml. of water to make approximately 0.48M 22 solution. The degassing bulb was charged with 250 ml. of the solution and the procedure repeated as before. Estimation of water collected in the trap from evaporation, gave the solution strength as 0.5MT 27e. ( i i ) Oxygen Concentration Method: Two oxygen concentra-tions were employed throughout the experiment, one was under one atmosphere of air and the other under one atmosphere of oxygen. For any particular temperature, the degassed sample tube plus two others of the type shown in Figure II (d) were placed in the circulating water bath, and the thermoregulator adjusted to the desired temperature. Four mis. of either d i s t i l l e d water or salt solution were placed in the sample tube. A polyethylene tube of 0.2 cm. O.D. connected to the oxygen cylinder was inserted through section A to the bottom of the sample tube. In the same manner, another polyethylene tube was connected to the air cylinder and inserted in another sample tube. Oxygen, or a i r , was bubbled slowly with stopcock A opened for about 15 minutes. At the end, each stopcock in turn was closed, the polyethylene tube withdrawn carefully, and the opening at A corked. The tube was dried on the outside and placed quickly in the induction c o i l for relaxation time measurements. Oxygen concentrations were checked occasionally by the micro-gasometric technique of Scholander et a l . (1955) and the results were - 2.5% of the theoretical concentration at that particular temperature. This is of the same order of magnitude as the usual error in the estimation of oxygen concentration by this method. ( i i i ) Measurement of Spin-Lattice Relaxation Time T^: In the process of having the three sample tubes equilibrating at the desired temperature and pressure, a glycerine sample tube was placed in the induction c o i l , and the Larmor frequency of the proton was adjusted to an optimum height signal. Then, the glycerine tube was replaced by a sample tube, and T^ determined for that particular sample. T^ measurements reported here were carried out using a slight modofication of the free induction technique pulse method as developed by Hahn (1950). At the Larmor frequency of the protons, transitions between the nuclear Zeeman levels by an r - f magnetic f i e l d 2H^ cos co t are applied perpendicular to the polarizing magnetic f i e l d H q in the form of pulses. H q in this experiment was supplied by a permanent magnet. The effect of the pulses i s to rotate the nuclear magnetization vector and thus establish components of the macroscopic nuclear magnetization in the plane perpendicular to the polarizing magnetic f i e l d . An alternating voltage i s thus set up by magnetic induction in the same c o i l as is used to produce ar-f magnetic f i e l d , and this voltage is amplified, detected and applied to the y -axis of an oscilloscope, the sweep of which has been triggered j u s t before application of the pulse. In using a TT/2- TT/2 pulse sequence the f i r s t pulse makes the magnetization wholly transverse (Mz=0), and the second pulse, applied at a time T a f t e r the f i r s t pulse, i s a search-ing pulse, the amplitude A ( T ) of i t s induction t a i l being a measure of the recovery of the z component of magnetization. However, i n such cases as relaxation time of water i s long and of the order of seconds and the induction t a i l amplitude of the second pulse i s nearly zero. A modification of the usual pulse technique was then used where the amplitude of the f i r s t pulse signal or the echo, which appears at twice the time of separation of the two pulses, could be employed to determine T^. A succession or t r a i n of pulses was applied. This t r a i n ensured that the magnetization a£ the end of the t r a i n was close to zero. At a time t, later, a s i m i l a r t r a i n of pulses was applied and the height A(t) of the f i r s t pulse signal was measured. T^ was deduced from the usual r e l a t i o n A(t) = A(oo) [ l-exp(-t/T 1) 1 (9 ) where A(*o) i s the value of A(t) for t » T^. Usually t was of the order of 4T^ before A(<*>) was obtained. Time was measured by a stopwatch; the t r a i n of pulses was operated manually by switch S (Figure II (a) ). Figure I I I shows the oscilloscope display of the two pulses and the detected r - f signal. The induction t a i l of the f i r s t pulse is shown at A ( ~ o ) and A ( t ) . Figure I V shows evaluation of T^ from a plot of log A A U A = A ( o o ) -A(t)"} as a function of time. 2 6 RESULTS AND DISCUSSION I. (i) Determination of T^, the spin la t t i c e relaxation time, for d i s t i l l e d water and 0.5M sodium chloride solution at various temperatures and different concentrations of oxygen. T^ values for oxygen-free and for oxygen-containing water under one atmosphere pressure of oxygen and one atmosphere pressure of air were obtained between 1°C. and 75°C. for d i s t i l l e d water and 1°C. to 40°C. for salt solution (Table I ) . The values were plotted versus temperature and the best lines were drawn by eye through the points as shown for oxygen-free d i s t i l l e d water in Figure VII. The lines were extrapolated to 0°C. and the smoothed values of T^ were taken at 5C° inter-vals and l i s t e d in Table II for double d i s t i l l e d water and in Table III for 0.5M sodium chloride solution. Figures V and VI show the best curves obtained for T^. The estimated maximum error for T^ data for the smoothed + curves was not more than - 5%. The results of T^ for oxygen-free d i s t i l l e d water were compared to those of Simpson and Carr (1958) who measured T^ and the diffusion coefficient D of proton in water at tempera-tures between 0°C. and 100°C. The results are shown in Figure VII. Agreement at temperatures above 40°C. was very apparent where the difference between the two results was of the order TABLE I T L (MEASURED) Di s t i l l e d Water Temperature Oxygen-free 1 Atm. Air 1 Atm. 0 2 °C. T x sec. T1 sec. T^ sec. 1.0 2.5, 2.6 2.0, 2.1 1.1, 1.2 7.0 2.9 2.5 1.4 9.0 3.0 2.4 1.5 12.0 3.2 2.8 1.8 15.0 3.2 - -17.0 3.5 3.0 1.9 21.0 3.9 3.2 2.3 25.0 3.9, 4.0 3.6 2.4 28.0 4.2 4.0 2.8 32.0 4.6 4.2 2.9 35.0 4.6 - -36.0 4.6 4.3 3.0 40.0 4.6, 4.7 4.5 3.3 44.0 5.4 4.8 3.6 48.0 5.5 5.2 3.7 52.0 6.2, 5.9 5.3, 5.6 4.1, 4.2 55.0 6.4 5.5, 5.8 4.1, 4.2 58.0 6.4, 6.8 5.9, 6.2 4.3, 4.5 64.0 7.1, 7.5 6.4, 6.9 4.8, 4.9 70.0 7.7, 8.0 7.2, 7.5 5.3, 5.5 75.0 8.3, 8.6 7.6, 7.9 5.8, 6.1 0.5M Sodium Chloride Solution 1.0 2.6, 2.7 2.0, 2.1 1.2, 1.3 8.0 3.0 2.7 1.8 13.0 3.3, 3.5 2.8 1.7, 1.9 15.0 3.4 - -18.0 3.7 3.0, 3.2 1.9, 2.0 23.0 4.0, 4.2 3.6 2.2 28.0 4.4, 4.6 3.8, 3.9 2.2, 2.4 32.0 5.0 4.1, 4.2 2.5 35.0 5.0, 5.3, 5.2 4.1, 4.4 2.7, 2.8 40.0 5.7, 6.0 5.1, 5.2 3.1, 3.3 TABLE II VALUES OF T AT DIFFERENT TEMPERATURES AND CONCENTRATIONS OF OXYGEN AS GIVEN BY SMOOTHED CURVE* Di s t i l l e d Water Temperature Oxygen-free 1 Atm. Air 1 Atm. 0 °C. T^ sec. T^ sec. T^ sec. 0 2.53 2.03 1.13 5 2.80 2.31 1.39 10 3.08 2.62 1.67 15 3.37 2.90 1.95 20 3.64 3.19 2.20 25 4.00 3.54 2.48 30 4.30 3.85 2.77 35 4.63 4.17 3.02 40 5.00 4.50 3.28 45 5.38 4.73 3.58 50 5.80 5.29 3.86 55 6.30 5.70 4.17 60 6.85 6.15 4.50 65 7.34 6.76 4.90 70 7.82 7.08 5.40 75 8.40 7.84 8.10 of 1-2%. At lower temperatures the difference was more pro-nounced, with Simpson and Carr's T^ being much lower. T^'s differed by 19% at 20°C. and 39% at 0°C. However, better agreement at lower temperatures was obtained with Guilloto's (1956) work, where the difference was only 4% at 20°C. * The values of Table. II were obtained at 5C° intervals from the smoothed curves of the plots of measured T^. These were employed in a l l subsequent calculations. 29 TABLE III VALUES OF T AT DIFFERENT TEMPERATURES AND * CONCENTRATIONS OF OXYGEN AS GIVEN BY SMOOTHED CURVE 0.5M Sodium Chloride Solution Temperature Oxygen-free 1 Atm. Air 1 Atm. 0„ °C. T^ sec. T^ sec. T1 sec. 0 2.62 2.10 1.21 5 2.90 2.40 1.43 10 3.20 2.67 1.63 15 3.50 2.94 1.82 20 3.81 3.24 2.04 25 4.20 3.58 2.27 30 4.70 3.98 2.53 35 5.28 4.47 2.80 40 6.00 5.05 3.21 The iron impurities in salt solution were of the order 15 -3 of 10 ions cm . Hausser and Laukien (1959) measured the 1 1-1 spin relaxation time probability y ^ for solutions of Fe ^ ^  I and Fe in aqueous solutions at f i e l d strength of 6200 -H- 21 -3 gauss. For concentrations of Fe of 10 ions cm T.., N 1(c) -2 was of the order of 10 sec. with l i t t l e variation with •••i-i 19 temperature. For concentrations of Fe of 2.6x10 ions -3 -2 -2 cm , T n, » varied between 10 to 3.5x10 sec. at temperatures 1(c) between 0°C. and 90°C. For such concentrations present as impurities in this experiment, the effect of paramagnetic * The values of Table III were obtained at 5C° intervals from the smoothed curves of the plots of measured T^. These were employed in a l l subsequent calculations. 30 ions i s assumed to be less than the experimental uncertainty of 5% i f the iron was in the ferrous state and slightly larger i f i t was totally in the ferric state. ( i i ) Behaviour of 1/T^ with oxygen concentration at different temperatures. The method used to produce two different concentrations of oxygen was simple and accurate to a few per cents of the theoretical values at a l l experimental temperatures. The only disadvantage was the incapability of attaining constant concentrations at different temperatures which was essential in the observation of oxygen effect on the relaxation time T^. It was then the primary concern in data presentation to show the variation of T^ with concentration and temperature. Tables IV and V and Figures VIII and IX show the behaviour of 1/T^ with oxygen concentrations both in d i s t i l l e d water and 0.5M sodium chloride solution. The graphs present clearly the linear relation of oxygen concentration to TABLE IV VALUES OF 1/T1 AT DIFFERENT OXYGEN CONCENTRATIONS AND DIFFERENT TEMPERATURES - DISTILLED WATER Oxygen free 1 Atm. Air 1 Atm. Oxygen Temperature Oo Cone. 0 2 Cone. 0 2 Cone °C. 1/TX mg./l. 1/TX mg./l. 1/TX mg./l. 0 .395 — .492 14.56 .884 69.45 5 .357 - .432 12.72 .719 60.72 10 .324 - .382 11.25 .598 53.68 15 .296 - .344 10.06 .512 48.02 20 .274 - .313 9.10 .454 43.39 25 .250 - .282 8.24 .403 39.31 30 .232 - .260 7.52 .361 35.88 35 .215 - .240 6.95 .331 33.15 40 .200 - .222 6.46 .304 30.82 45 .185 - .211 5.99 .279 28.58 50 .172 - .189 5.57 .259 26.57 55 .158 - .175 5.12 .239 24.45 60 .146 - .162 4.77 .222 22.74 65 .130 - .148 4.20 .204 20.03 70 .127 - .141 3.89 .185 18.56 75 .119 - .127 3.38 .164 16.10 TABLE V VALUES OF l/T AT DIFFERENT OXYGEN CONCENTRATIONS AND DIFFERENT TEMPERATURES - 0.5M SODIUM CHLORIDE Oxygen free 1 > Temperature 0_ Cone. °C. l/T± mg./l. UT1 0 .381 .476 5 .344 .416 10 .312 .374 15 .285 .340 20 .262 .308 25 .238 .279 30 .212 .251 35 .189 .223 40 .166 .198 Air 1 Atm. Oxygen O Cone. O2 Cone mg./l. 1/TX mg./l. 11.64 .826 55.52 10.20 .699 48.65 9.05 .613 43.17 8.14 .549 38.82 7.41 .490 35.35 6.81 .440 32.48 6.30 .395 30.05 5.83 .357 27.81 5.35 .311 25.52 32 To examine for the linearity, an added test was performed before the f i n a l plot was completed. From the best drawn lines for T^ of oxygen-free water and T^ for oxygen-concentrated water under 1 atmosphere of oxygen, 1/T-^  values were obtained and plotted versus oxygen concentrations for different tempera-tures. A straight line was drawn between the two values at each temperature. For oxygen concentration under one atmosphere of a i r , 1/T^ was read at the concentration of oxygen at that particular temperature and compared with that of 1/T^ obtained from the experiment; the two values were within ^ 1.07o of the smoothed values for a l l temperatures. Although the individual measurements of T^ were only accurate to about presumably the drawing of a smooth curve through the plots of T^ versus temperature has the effect of averaging out the random errors to some extent, Another characteristic in Figures VIII and IX is revealed by the change of line slope with change of temperature. This now becomes the most significant aspect of the experiment, for though the linearity of 1/T^ with concentration is maintained, the effect of dissolved oxygen on relaxation times is shown to vary as a function of temperature. In addition, a constant concentration line can be obtained from the linearity relation. This w i l l be a measure of the relaxation time effect under the same concentration. A vertical line at 32.00 mg. O2/I. was 33 drawn through the slope lines, and the slopesci determined and plotted as a function of temperature in Figures X and XI. The values corrected for the best line on the graph are given in Tables VI and VII. The lines represent 1/T^^ at constant 2/ -3 concentration of of 10 moles/1, where 1/T-^^ = ^/ Tl (o ) 1 / T l ( p u r e ) * TABLE VI 1/T1 VALUES OF OXYGEN-FREE AND 32.00 MG./L. OXYGEN-CONCENTRATED-DISTILLED WATER corrected Temp. °C. 1 / T Ko 2 ) l/T l(pure) 1 / T K c ) 1 / T l ( c ) co: 0 0.607 0.395 0.211 0.211 5 0.547 0.357 0.190 0.186 10 0.487 0.324 0.162 0.162 15 0.440 0.296 0.143 0.143 20 0.407 0.274 0.132 0.132 25 0.374 0.250 0.124 0.123 30 0.348 0.232 0.115 0.116 35 0.326 0.215 0.110 0.110 40 0.307 0.200 0.107 0.107 45 0.290 0.185 0.104 0.104 50 0.276 0.172 0.103 0.103 55 0.263 0.158 0.104 0.103 60 0.250 0.145 0.104 0.102 65 0.239 0.136 0.102 0.102 70 0.231 0.127 0.103 0.102 75 0.210 0.119 0.091 0.102 34 TABLE VII 1/T VALUES FOR OXYGEN-FREE AND 32.00 MG./L. OXYGEN-CONCENTRATED-0.5M SODIUM CHLORIDE SOLUTION Temp. °C. l/Ti (o 2 ) 1/Tl(pnre) 1 / T l ( c ) 1 / T ( c ) corrected 0 0.639 0.381 0.257 0.257 5 0.576 0.344 0.231 0.234 10 0.533 0.312 0.221 0.221 15 0.502 0.285 0.216 0.212 20 0.468 0.262 0.205 0.205 25 0.434 0.238 0.196 0.199 30 0.406 0.212 0.193 0.195 35 0.382 0.189 0.192 0.192 40 0.357 0.166 0.190 0.190 Concentration values of oxygen in 0.5M sodium chloride solu-tion were obtained from the published results of Truesdale et. a l . (1955) for 307oeS. The values of d i s t i l l e d water, on the other hand, were taken from Handbook of Physics and Chemistry (1960), p. 1706. The use of two different data sources was merely a mea-sure of convenience, since the results of Truesdale et a l . were limited to temperatures not exceeding 40°C. The most apparent effect of oxygen on the proton relaxa-tion time is shown by the curve of 1/T^^C^ as a function of temperature for the same concentration of molecular oxygen. T,/ v is seen to be larger for d i s t i l l e d water than for 1(c) & NaCl solution at temperatures between 0° and 40°C. If the change in relaxation time is to be considered solely as a function of viscosity or diffusion, then qualitatively the results w i l l be in agreement with the data of oxygen diffusion 35 in d i s t i l l e d and salt solutions. Diffusion coefficient, D, values for molecular oxygen were obtained by Rotthauwe (1958). By means of a dropping mercury electrode he found the diffusion coefficients of 0 2 in air-saturated pure and sea water at 20°C. to be 2.92xl0"5 cm2 sec." 1 for d i s t i l l e d water, 2.8xl0" 5 2 -1 -5 2 -1 cm sec. for 16%o salinity seawater and 2.72x10 cm sec. for 35%o salinity seawater. However, as is shown in the next section, viscosity and molecular diffusion play a part in the relaxation mechanism, and although the part played is a major one, other aspects of complex formation enter the picture and contribute to the decrease of l / T ^ ^ with tempera-ture. II. Relaxation Terms If the relaxation time effect T^( c) for the special case of 0 2 in water depends on two mechanisms of relaxation, as was assumed in the theoretical part, where one is a dipole-dipole interaction and the other an exchange or hyperfine interaction as in ^ = cTc+wlV (8), then, the dependency of l / l i ( c ) on the precession frequency of the electron and hence on the strength of the magnetic f i e l d H Q becomes evident. The form of equation (8) is obtained by the assumptions placed on the correlation times v c and s e 36 whose validity depends on the proper evaluation of the different terms and on the understanding of oxygen complex formations. In such a case, the three essential correlation times Tg» *r a n a T^ which are: the relaxation time of the electron, the tumbling relaxation time, and the lifetime of the complex respectively, should be estimated approximately to allow for the conditions where T Ldc«l and ^ 0j o» 1. c " e " The dependence of T^ in part on the dipolar interaction and hence on temperature is attributed to variation in where -rr— = — i — + —^—: thus 7^" w i l l be a function of the ^ c T r nrs tumbling of complex only i f 7" is much smaller than rr. • In the light of experiments done on paramagnetic ions, the relaxation correlation time of the electron Tg was found to _q _ i (] i i J_4- 4 4 | i i be between 10 to 10 sec. v for Mn , V , Cu , C r m and Ni (Bloembergen and Morgan, 1961). On the other hand, *T and consequently ff" become the correlation time for the r c tumbling of the water oxygen complex. Since the amount of dissolved oxygen is of the order of 10 moles/1., then pro-bably not more than one oxygen molecule forms a complex with a water cluster at any one time. The presence of oxygen should be of no consequence in regard to viscosity, and hence would allow for the use of viscosity correlation time as related to the tumbling of the cluster. Thus T temperature dependence (Cox and Morgan, 1959), T = T ° e E c/ R 1- becomes 37 that of the water or the salt solution in which the oxygen is dissolved. However, the exponential dependence does not necessarily follow for the spin exchange contribution, nor for the dipolar contribution when is not primarily determined by the tumbling of the complex, but is also determined by the time dependence of the orientation of the electron spin v g. Similarly the argument holds for the exchange interaction where _ i _ = _L_ + _L_ # Both a n (* ^ h a r e temperature dependent, ^ e 1 S 1 h with increasing with temperature and magnetic f i e l d and ^ k decreasing with temperature only. For paramagnetic ions in general (Bloembergen, 1957; Bloembergen and Morgan, 1961), the relation between the three correlation times is of this order: ^ ^~g "^^x* This order may be similar for oxygen except that in the latter case ^ g may be greater than ^ j j , the lifetime of the complex. This follows from the reasoning that the oxygen molecule is non-polar and thus incapable of forming a hydration layer of the same strength or magnitude as that of a charged ion in solution. Some measure of the relative magnitudes of the three time constants discussed above can be derived from the dif-ference between values of l/T-j^ c^ at the same temperature and concentration of oxygen but at different magnetic f i e l d strength. The work of Guilloto (1956) which was done at a f i e l d strength 38 o of 1650 gauss and at 20 C. provided a point of comparison at that temperature as well as means of evaluating equation (8). Furthermore, the use of his results was augmented by the close agreement of for degassed d i s t i l l e d water reported by him and which differed only by 4% from the value reported here. His value for 1/T^^ at oxygen concentration of 6x10^ -3 -1 -1 molecules cm was 0.354 sec. in comparison to 0.132 sec. obtained in this experiment. It is impossible to f i t a l l the f i e l d dependent terms uniquely in equation (3) with values of 1/T^^ at only two different fields at one temperature. However, i t may be f r u i t f u l to consider two limiting cases. Case I If the hyperfine interaction term was assumed to be 2 negligible, i.e. < A > was equal to zero, and equation (3) solved from the values of l/T.. x at the two frequencies, no 1(c) solution is possible for the 1/T^^ ratios of 2.67 obtained from Guilloto and this work. The largest value that (l/T.,, S)H = 1650/(1/T W J H = 7000 can have for < A2> =0 1(c) o 1(c) o is 2.1 and this maximum is reached at T s 3x10 sec. c One therefore has to assume that an error of at least 30% exists in the ratio of the two results to be able to solve the equation where the f i r s t near ratio to 2.67 occurs for T = 3 x l 0 _ 1 1 sec. at 2.1 c 39 If T , the correlation time of water, stands for the c tumbling of the complex T , then two d i f f i c u l t i e s present themselves when the above approach is considered. The f i r s t is due to the value of T c for water which is usually of -19 the order of 10 sec. From data of dielectric relaxation (Eigen and Maeyer, 1958), neutron diffraction (Brockhouse, 1958), diffusion and viscosity (Wang ejt a j L ., 1953) , i t is seen that the largest correlation time T for pure water at 0°C. i s -11 about 10 sec. The model representing this correlation time is the well known equation of Debye: T - do) Bloembergen (1961) related T" to T c by the equation T = 3 T saying that Debye's ^ refers only to the orientation of the polar group in space, while for T the relative orienta-tion between the magnetic nuclei must be considered. However, the characteristic time ^" of Debye and the correlation time in the magnetic local f i e l d spectrum are proportional in one sample. They both vary in proportion to *?/T i f the temperature of the water is changed. This leads into the second d i f f i c u l t y where, above the temperature of 45°C, 1/Ti( c) becomes nearly constant. Thus i f only the f i r s t term of relaxation i s considered, i t means that is unchanged between 45° to 75°C. and this is very unlikely. 40 Further, i f the assumption is made that the anomaly discussed above is due to ^"g» t n e relaxation time of the electron, then correlation with the experimental results at higher temperatures proves d i f f i c u l t . Bernheim et a l , (1959) found that Tg for paramagnetic ions changes linearly with temperature in the fashion T " T o % increasing with increase in T ° rr temperature. The increase of l g with temperature should further remove i t from influencing f^"c, and thus makes ^ c more dependent on T f , the tumbling time of the complex. Case II When tT"c value is assumed to be sufficiently below 1 0 - 1 1 sec. over the entire temperature range that ( C J g T c ) 2 «• 1 and (CJ T ) 2 ^ 1, equation (8) may be used to interpret the o e results. S e The assumption that £J) s T e » 1 w i l l be dis cussed in terms of the behaviour of 1/T^ C^ curve at higher temperature, showing f i r s t that < A > is not zero, and secondly CO r T is much o e larger than one. The coefficients C T c and nj^ can be evaluated at 20°C. from the two equations at 6J g equals to 1.24X1011 sec. 1 used in this work and CO g equals to 2.92x10^ secT 1 for the results of Guilloto. C T_ and — - were obtained as 10.40xl0~2 sec! 1 - cu 2 T S e respectively at the frequency O) of 1.24x10*''*" secT*" the value appropriate to a l l the temperature dependent measure-ments of this experiment. r-r B (i) Evaluation of C T and — n . at different tempera-c 60? T o e tures for d i s t i l l e d water and 0.5M sodium chloride solution. In order to get information on T"e as a function of tempera-ture, an assumption has been made about the temperature depen-dence of T . The assumption was based on two considerations: 1. That oxygen in solution does not alter the size of the complex appreciably nor the water viscosity, and 2. That T is dependent mainly on the tumbling time T of the complex. Guided by other workers (Simpson and Carr, 1958; Bloembergen, 1961, pp.84-91), c T c was considered a function of ^/T, and was evaluated at a l l experimental temperatures. Tables VIII and IX give the values for C T obtained c from viscosity (Int. C r i t . Tables, 1923; Dorsey, 1940, p. 188) and absolute temperature data for both d i s t i l l e d water and salt solution at a constant concentration of 10 moles/1, of oxygen. The results are also shown graphically in Figure XII. The term (C T ) was calculated for a l l temperatures by starting with the determined value from the two frequencies at 20°C., using the following relationship. n (C T )Ta - a / T a (C T )Tb (11) C ^b/Tb C 42 TABLE VIII EVALUATION OF C T VALUES AT c . DIFFERENT TEMPERATURES DISTILLED WATER Temperature tyTxlO3 1 rr- ? °C. c.p. c.p. deg" 1 c T x i o z C 0 1.7921 6.564 20.00 5 1.5188 5.463 16.65 10 1.3077 4.620 14.08 15 1.1404 3.960 12.06 20 1.0050 3.430 10.40 25 0.8937 3.000 9.14 30 0.8007 2.642 8.05 35 0.7225 2.346 7.15 40 0.6560 2.095 6.38 45 0.5988 1.883 5.74 50 ' 0.5494 1.700 5.18 55 0.5064 1.544 4.70 60 0.4688 1.407 4.29 65 0.4355 1.288 3.92 70 0.4061 1.183 3.60 75 0.3799 1.091 3.32 TABLE IX EVALUATION OF C T VALUES AT DIFFERENT TEMPERATURES -c 0.5M SODIUM CHLORIDE SOLUTION Temperature ^/TxlO °C. c.p. c.p. deg 0 1.8727 6.859 5 1.5978 5.747 10 1.3822 4.884 15 1.2065 4.189 20 1.0703 3.653 25 0.9553 3.206 30 0.8632 2.849 35 0.7740 2.513 40 0.7010 2.240 " 1 C TxlO 2 secT1 c 20.90 17.51 14.88 12.76 11.08 9.77 8.68 7.66 6.82 43 B The term — _ was evaluated at d i f f e r e n t temperatures ^ 5 e from the difference of 1/T,, N and C T and given i n Tables X 1(c) c ° •p and XI. Figure XIII shows the v a r i a t i o n of • • w ^ with tempera-ture, where for d i s t i l l e d water Q"e decreases l i n e a r l y , and for the s a l t solution exponentially. I t i s also seen that i n s a l t solution i s smaller at a l l temperatures and decreases fas t e r for the same concentration of oxygen. Behaviour of thi s nature i s consistent with the assumption that CO c *T >> 1; the contribution of the second term to 1/T-j^^ would not increase with increase of temperature i f was either nearly equal to one or less than one. Also, the decrease of T g with temperature shows the dependency i n t h i s respect to be more on than on T , since T i s also increasing with h b o temperature. ( i i ) Comparison of A c t i v a t i o n Energies and Heats of Solution. By i d e n t i f y i n g the dependence of T on the l i f e t i m e of the complex from the above consideration, an exponential temperature dependence may be assigned to T e because i t s v a r i a t i o n now i s only a function of temperature, independent rr- rr- (o) -Ee/RT of the magnetic f i e l d . The re l a t i o n s h i p T e = T e e would allow for the determination of the a c t i v a t i o n energy . . ( 0) -EC/RT E , i n the same manner as i n T = T e . The difference e' v c v c between the two a c t i v a t i o n energies l i e s e s s e n t i a l l y i n the 44 concept treatment of the complex. For the activation c energy w i l l be solely a measure of hydrogen bond breaking between water molecules (Wang et a l . , 1953), due to the tumb-ling of the complex, disregarding the size contribution of oxygen. On the other hand, E £ i s more a measure of the bond broken between the oxygen molecule and the cluster forming the complex. If the heat of solution of oxygen in water, - xv H is actually an indication of the bond formed between the oxygen molecule and water, then, the comparison of - &• H and E should show a simultaneous change under various conditions. TABLE X VALUES OF 5 — AND - J - AT DIFFERENT TEMPERATURES -DISTILLED WATER Temp. °C. T exlO~ sec. -20 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75 21.18 18.63 16.24 14.33 13.22 12.33 11.61 11.06 10.71 10.45 10.36 10.32 10.29 10.27 10.25 10.23 20.00 16.65 14.08 12.06 10.40 9.14 8.05 7.15 6.38 5.74 5.18 4.70 4.29 3.92 3.60 3.32 1.18 1.98 2.16 2.27 2.82 3.19 3.56 3.91 4.33 4.71 5.18 5.62 6.00 6.35 6.65 6.91 1.82 3.05 3.33 3.50 4.34 4.91 5.48 6.02 6.67 7.25 7.98 8.65 9.24 9.78 10.24 10.64 45 TABLE XI VALUES OF ^2 AND T e AT DIFFERENT TEMPERATURES -S e 0.5M SODIUM CHLORIDE SOLUTION 1/T 1 ( c ) C T B B Temp. 9 L { C ) . . C _ Z 3 1 p T 7% °C. xlO sec" xlO sec. x l 0 2 s e c l l xl0" 2 0sec. 0 25.74 20.90 4.84 7.45 5 23.47 17.51 5.96 9.18 10 22.18 14.88 7.30 11.24 15 21.22 12.76 8.46 13.03 20 20.50 11.08 9.42 14.51 25 19.92 9.77 10.15 15.63 30 19.50 8.68 10.82 16.60 35 19.22 7.66 11.56 17.80 40 19.03 6.82 12.21 18.80 The activation energy E c was calculated in Figure XIV on the basis of data obtained from Tables XII and XIII. E is shown to be 4.4 kcal. for d i s t i l l e d water and 4^9 kcal. c for 0.5M sodium chloride solution. E was calculated from e Tables XIV and XV and Figures XV and XVI, using the relation rr = rr (o) e~Ee/RT e e The heat of solution for oxygen solubility in d i s t i l l e d water and salt solution was obtained by using the constant K of equation (1): P^ = KC^, where P^ = x(P-p w); x = frac-tional content of 0 2 in atmospheres, P = atmospheric pressure; and p w = water vapor pressure in the van't Hoff equation for enthalpy. The term K was determined at different temperatures under one atmosphere of oxygen as shown in Tables XVI and XVII, 46 and the results presented graphically in Figure XVII. Both the plots for E and - A H are seen to be non-linear e and hence are temperature dependent. Evaluation of E £ and - * H at temperatures of 0°, 15°, 30° and 60°C. showed a decrease with temperature in the same trend and of the same order of magnitude, thus revealing the close association between the two quantities (see Table XVIII). In general, both E and - A H are smaller for salt solution than for d i s t i l l e d water, and their decrease with rise in temperature proceeds at a faster rate. This could be a t t r i -buted to the weaker nature of the water-oxygen complex in the salt solution and to a higher dependence on temperature change. There is thus an indication that the cluster size determines the strength of the bond formed in the complex between oxygen and water, and hence the solubility. The anomaly at 0° to 15°C. in E g could be interpreted as a result of subtracting two large terms to obtain a small term. A small percentage error in C T"c or 1/T^^ at these temperatures B , w i l l appear highly magnified in 2 rr • Also the use of four 0 ) s <e or five significant figures in data tables such as those of vi s -cosity and logarithm values, was simply a matter of straight quotation rather than an indication of the expected accuracy. 47 TABLE XII ESTIMATION OF ACTIVATION ENERGY OF T -c DISTILLED WATER .CT = C T ° e " E / R T c c >erature °C. 1000 , -1 T cleg K r0/ r c c 2.303 log T°/ c 0 3.650 5 3.597 1.201 0.1831 10 3.533 1.420 0.3507 15 3.472 1.658 0.5057 20 3.412 1.923 0.6540 25 3.355 2.188 0.7833 30 3.300 2.484 0.9101 35 3.246 2.797 1.0287 40 3.194 3.135 1.1427 45 3.144 3.484 1.2485 50 3.095 3.861 1.3512 55 3.048 4.255 1.4484 60 3.003 4.662 1.5398 65 2.958 5.102 1.6301 70 2.915 5.556 1.7153 75 2.873 6.024 1.7961 E = 4.4 kcal. The contribution of the two relaxation terms C T~ and _ B (jj 2 fr is of importance in correlating the experimental results S e with oxygen solubility in the two solvents. The effective contribution of the term C T is a function of the viscosity and/or diffusion, and the contribution of the second term •p —K=— a function of the oxygen water complex formation. The ratiot of the two terms at various temperatures is obtained from Tables XIV and XV and plotted in Figure XVIII. 48 TABLE XIII ESTIMATION OF ACTIVATION ENERGY T -c 0.5M SODIUM CHLORIDE SOLUTION Temperature 1000, -1 „ T °/ T 2.303 log T °/ f r 0 3.650 5 3.597 1.194 0.1773 10 3.533 1.405 0.3399 15 3.472 1.638 0.4938 20 3.412 1.886 0.6347 25 3.355 2.139 0.7605 30 3.300 2.408 0.8788 35 3.246 2.728 1.0039 40 3.174 3.065 1.1202 E = 4.9 kcal. A striking difference appears between the two ratios when the two terms become equal. For salt solution the equality occurs at 30°C, while for d i s t i l l e d water, i t takes place at 50°C. Because of l i t t l e difference in the terms c (Figure XII) between d i s t i l l e d water and salt solution, the anomaly seems to be due largely to ^ and i t s dependence on ion content in water, or more correctly on the nature of the cluster formation. The breaking of such water aggregates by either temperature or ions, clearly shows in the variation of T . The close association in this case between cluster e and complex formation renders as a valuable parameter in the evaluation of liquid properties and the extent and nature of bonding of the dissolved substance. 49 TABLE XIV B RATIO OF C T //,\ 2 rr AND ESTIMATION OF ACTIVATION ENERGY OF T -c g v e e DISTILLED WATER Temp. 1000 , -1„ o c^ —rjT- deg ..K B Ratio C t / ^ t r ^efo) 2.303 log TeCo) 0 3.650 16.95 5 3.597 8.41 1.678 0.5177 10 3.533 6.52 1.830 0.6045 15 3.472 5.31 1.924 0.6545 20 3.412 3.64 2.390 0.8715 25 3.355 2.87 2.703 0.9947 30 3.300 2.26 3.017 1.1045 35 3.246 1.83 3.314 1.1983 40 3.194 1.47 3.669 1.3000 45 3.144 1.22 3.992 1.3846 50 3.095 1.00 4.390 1.4797 55 3.048 0.84 4.763 1.5612 60 3.003 0.72 5.085 1.6266 65 2.958 0.62 5.381 1.6833 70 2.915 0.54 5.636 1.7296 75 2.873 0.48 5.856 1.7678 TABLE XV B RATIO OF c T / ^ f r AND ESTIMATION OF ACTIVATION ENERGY OF T -c g e e 0.5M SODIUM CHLORIDE SOLUTION B Temp. 1000 -1 Ratio C T 2 fr Te(o) 2.303 log ^ ( p ) o C. x deg K c ^ g ^e ^ 0 3.650 4.32 5 3.597 2.94 1.232 0.2087 10 3.533 2.04 1.509 0.4115 15 3.472 1.51 1.749 0.5592 20 3.412 1.18 1.948 0.6669 25 3.355 0.96 2.098 0.7411 30 3.300 0.80 2.236 0.8049 35 3.246 0.66 2.389 0.8712 40 3.194 0.56 2.523 0.9256 50 TABLE XVI EVALUATION OF HEATS OF FORMATION FOR SOLUBILITY OF OXYGEN UNDER ONE ATMOSPHERIC PRESSURE OF Ort - DISTILLED WATER Water Vapor Q^:Cone. POo 2.303 log Temp. Pressure 0_ moles/1. u2 c 2 °C. mm.Hg. 1 xlO 3 Corr. Atm. P 0 2 0 4.6 2.170 0.993 -6.1269 5 6.5 1.897 0.991 -6.2596 10 9.2 1.677 0.988 -6.3800 15 12.8 1.500 0.983 -6.4811 20 17.5 1.355 0.978 -6.5829 25 23.8 1.228 0.968 -6.6711 30 31.8 1.121 0.958 -6.7582 35 42.2 1.035 0.945 -6.8180 40 55.3 0.963 0.928 -6.8719 45 71.9 0.893 0.905 -6.9221 50 92.5 0.830 0.879 -6.9663 55 118.0 0.764 0.845 -7.0099 60 149.4 0.710 0.804 -7.0334 65 187.5 0.625 0.754 -7.0967 70 233.7 0.580 0.692 -7.0856 75 289.1 0.503 0.620 -7.1181 TABLE XVII EVALUATION OF HEATS OF FORMATION FOR OXYGEN SOLUBILITY UNDER ONE ATMOSPHERIC PRESSURE OF 0„ - 0.5M SODIUM CHLORIDE SOLUTION Temp. °C. Water Vapor Pressure mm.Hg. C^:Cone. O0 moles/1. xlO 3 P o 2 Corr. Atm. 2.303 log 0 2 P 0 2 0 5 10 15 20 25 30 35 40 4.5 6.2 7.1 13.0 17.3 23.3 31.4 42.3 54.5 1.735 1.520 1.349 1.213 1.104 1.015 0.939 0.869 0.797 0.994 0.992 0.988 0.983 0.978 0.970 0.959 0.945 0.929 -6.3517 -6.4823 -6.5976 -6.6987 -6.7876 -6.8639 -6.9300 -6.9928 -7.0621 51 TABLE XVIII HEAT OF SOLUTION AND ACTIVATION ENERGY OF DISSOLVED OXYGEN IN WATER Temperature D i s t i l l e d Water 0.5M Sodium Chloride Solution °C. % - A H Eja - A H 0 4.4 5.2 6.2 4.8 15 4.4 3.4 3.6 2.9 30 3.6 2.7 2.4 2.4 60 2.8 1.5 52 GENERAL DISCUSSION The case of complex formation in an oxygen-water system is seen to be fundamentally related to hydrogen bond formation in the same manner as that of cluster formation of water molecules. This is evident in the behaviour of oxygen solubi-l i t y and i t s dependence on the two important parameters of temperature and ion content. The st a b i l i t y of the hydrogen bond in this aspect may be considered due to either the orientation of the solvent and solute molecules to each other or to the dependence on the size and shape of the water cluster or on both. The decrease in T e with higher temperatures, and additions of ions to water would allow for the increase in diffusion of oxygen molecules because of the more random orientations and the more frequent jumps of oxygen between the clusters. However, a comparison of diffusion coefficients for ion-free and ionized water at the same temperature (Rotthauwe, 1958), shows the diffusion as decreasing with higher salt content, thus making i t more a function of the viscosity than of the stab i l i t y of the complex. Conversely, a lowering of the viscosity by an increase of temperature in ion-free water increases the diffusion coefficient from 1.87 to 2.60x10 cm2 sec",1 between 16° and 25°C. (Millington, 1955; Kolthoff and Miller, 1941). The change of viscosity in ionized solu-tions is affected by the ions in the neighbourhood of water 53 molecules. Bernal and Fowler (1933) suggested that the ions form a lat t i c e structure with water molecules, and the struc-ture tightening or loosening is due to extent of hydration of the ions. These structures, however, would be different than the normal clusters formed by ion-free water, and their orientation at the edges are expected to be modified by the f i e l d of nearest charged ion. The model could be pictured then, that an oxygen molecule attaches i t s e l f to a water cluster for a time T , this being a function of temperature and ion content, then in short lat t i c e jumps i t hops from one cluster to the other within a very small regional environment but does not contribute effectively to diffusion. Thus i f each regional environment remains in the same space allotted to i t , diffusion then becomes so small as to be of the same order as diffusion in solids. In liquids essentially, the a b i l i t y to move and exchange regional environments by mass movements of molecules or clusters, becomes the prime factor in determining the magnitude of diffusion. The energy of activation, E , obtained from viscosity c and absolute temperature data for d i s t i l l e d water is 4.4 kcal./mole and for 0.5M sodium chloride solution is 4.9 kcal./ mole. This is a measure of the energy required for one hydrogen bond breaking between two water molecules (Wang et a l . , 1953). A significant characteristic of this quantity is i t s constancy with temperature. Many experiments were performed by different workers to calculate the enthalpy of a hydrogen bond in water. Their values were 3.4 kcal./mole (Scatchard et, a l . , 1952); 4.4 kcal./mole (Rowlinson, 1949), and 5.7 kcal./mole (Lambert, 1953) at temperatures between -65° to 100°C. Also the activa-tion energy for dielectric absorption in water was found to be 13.2 kcal./mole (Autry and Cole, 1952) and was interpreted as the breaking of three hydrogen bonds each being 4.5 kcal/mole The temperature independence of hydrogen bond breaking and forming between water molecules distinguishes i t from the process occurring between oxygen molecules and water molecules where temperature changes influence the energy of bond forming or breaking. At temperatures of 0°C. and 15°C. the activation energy E is 4.4 kcal./mole and - i H 5.2 and 3.4 kcal. (Table XVIII). e For the 0.5M sodium chloride solution E varies between 6.2 and e 3.6 kcal./mole and - A H between 4.8 and 2.9 kcal/mole. These values of hydrogen bonds formed and broken for an oxygen water complex are of the order of the hydrogen bond strength formed between one water molecule and another at any temperature The proximity of the values of E £ to -AH for an oxygen-water system supports the assumptions made in the treatment of the experimental results. This agreement with the values of activation energy and enthalpy for water-water system at 55 low temperatures (0° - 15°C.) shows that: 1. Either two bonds are formed between the oxygen molecule and water with each bond being of the order of 2.2 kcal./mole, or 2. only one bond exists between water and oxygen being of the same magnitude as that formed between two water molecules. The nature of the bond in an oxygen-water system must be different from that formed in a water-water system. The energy involved in an interaction resulting in a hydrogen bond in the case of water molecules is always the same whether the process occurs at -65°C. or in the vapor phase at 100°C, and i s only modified by the presence of ions in solution. Conversely, the decrease of E e and - AH of oxygen-water bond at higher temperatures indicates that the forces acting on an oxygen molecule, in hydrogen bond formation, are changing continuously. At such temperatures these could be related to the increase in molecular agitation and thus to more dis-tortion of their line of approach where interaction between the dipole moment of water and the induced polarization of the oxygen molecule tends to be less effective. An explanation may be put forward to account for the change in forces as being dependent on either the size and form of the water cluster, or the distance R of the 0-H 0 bond or the angle of approach. Hughes et a l . (1956) interpreted the weakening of a hydrogen bond as i t appears in the stretching frequency ^ q i - n infra-red 56 as the decrease in the force f i e l d on a molecule due to i t s neighbours. As the temperature rises, the intensity of the force f i e l d decreases owing to the greater average distance between molecules. Earlier Cogeshall and Saier (1951) ex-plained the decrease in p g in a system of carbon tetrachloride-benzyl alcohol with increase in temperature as being dependent on the following characteristics: 1. Hydrogen bonding systems involve monomeric and usually several polymeric species in rapid equilibrium. 2. Each polymer has a characteristic 0 G , and the higher the polymer the lower i s P . At higher temperatures v)g i s increased, and this is supposed to accompany the decrease in the size of the polymer. It could be deduced that the stronger hydrogen bonds of oxygen-water are formed in the larger clusters. The decrease of the cluster size by molecular agitation due to an increase in temperature or by the presence of ions lowers the resultant force acting on polarizability of an oxygen molecule. The effect due to the change in the distance R of the 0-H 0 bond is dealt with by Pimentel and McLellan (1960, p. 242). The enthalpy AH is shown as a function of R for different isolines of electrostatic energy of point charge dipoles. For 0.30 electron charges, AH varies by 1 kcal./mole for a change of 0. in the bond length R. On the other hand, 5 7 the angles of approach are significant in the variation of &H. Figure XIX shows two ways of approach for two dipolar molecules and the designation of the angles © and <p between them (Pimentel and McClellan, 1960, p. 242). Although (J) could not be involved in the formation of a hydrogen bond in the case of water-oxygen due to the non-polarity of the oxygen molecule, where the induced polariza-tion is not governed by position, O s t i l l could contribute to the decrease of - AH and E by virtue of orientation of e J the water dipole molecule whether be i t the resultant dipole moment, or simply one of the 0-H moments. 58 CONCLUSION The results of this work reveal the following new points: 1. The solubility of oxygen in water is related to hydrogen bond formation between oxygen and water molecules. 2. The increase and decrease of solubility depends on the factors governing the building or destroying of clusters of water molecules. 3. Temperature and salt content affect the molecular orientation and the average line of approach of the reactants. 4. The effect of temperature and salt content appears in the activation energy and enthalpy. This is related to bonds broken and bonds formed consecutively. 5. The values of E £ and - *H are a measure of solubility. SUGGESTIONS FOR FURTHER WORK 1. For supplementing the assumption made for the oxygen water-cluster complex and the dependence of i t s dipole-dipole interaction correlation time on the viscosity of water, a l l measurements may be repeated at another frequency and magnetic f i e l d strength, and equation (3) solved accordingly. 2. Various other salt concentrations, and different non-paramagnetic salts may be used to examine the change in heats of solution and activation energies. 3 . Another paramagnetic gas, n i t r i c oxide, which contains a permanent dipole moment, may be used in a similar manner and the different values for solubility and relaxation times compared to oxygen for further understanding of the nature of the complex. 60 PART II ASPECTS OF OXYGEN SUPERSATURATION IN SEA WATER INTRODUCTION In every well mixed water body exposed to and at eq u i l i -brium with the atmosphere, the concentration of oxygen at constant temperature depends on Henry's law: P 2 - K C 2 (1). Any deviation from this relationship which displaces the equilibrium, or results in i t s invalidity, indicates certain physical, chemical or biological processes in action. These processes and their intensities are the prime factors which implement oxygen under- or supersaturation. "Oxygen Supersaturation" as i t applies in this work and as i t is commonly referred to (Klots, 1961) in sea, lake, or laboratory system could be defined as follows: A certain quantity of oxygen gas dissolved per unit volume of water in excess of a known saturated value (e.g. that given by Richards and Crowin (1956) for the particular temperature and ion content of the water under one atmosphere of a i r ) . The change of concentration of oxygen in the sea indepen-dent of exchange with the atmosphere is governed by depletion 61 and production. The f i r s t may be caused by oxidation of metal ions of lower valency such as iron I and II and manga-nese II and III; decomposition of organic matter; and respira-tion of plants and animals. On the other hand, production results from the process of photosynthesis in green plants only, and is limited to the euphotic zone (Richards, 1957). In natural water bodies oxygen content in the superficial layers often strongly deviates during the whole period of biological activity to either side of 100% saturation. The non-equilibrium content of oxygen in the surface layers under certain conditions may be due to rapid changes of the air temperature near the surface or the mixing of two oxygen-saturated water masses at different temperatures. Also a spe-c i a l but not experimentally demonstrated oxygen supersaturation may be produced at great depths in the ocean i f saturated water at the surface sinks at constant temperature and sal i n i t y . This results from the increase of Henry's law constant with depth. The reciprocal 1/K varies between 1.00 at the surface to 0.88 under 100 atmospheres pressure at 25°C. (Klots, 1961). A l l these factors cause but a short-lived deflection from equilibrium. However, a prolonged supersaturation of the surface water layers with oxygen points to the predominance within the water basin of processes of synthesis of organic matter (primary production) over processes of their destruction by respiration and oxidation (Vinberg, 1940). 62 Occurrences of this nature are frequent in the sea, and are more pronounced and varied in the neritic province and at higher latitudes than in the oceanic province and lower latitudes. The influence of land masses in providing more organic and inorganic matter as nutrients, and the more extreme variation in weather conditions of the higher latitudes on these regions, subject them to exhibit characteristics widely different from those of the more homogeneous nature that take place far at sea or in inshore tropical regions. The proximity of land is also believed to render the coastal waters as a whole about f i f t y times more productive in phytoplankton than the open ocean waters (Sverdrup, Johnson and Fleming, 1942). Reported oxygen data for several years in the waters of the West Coast of Canada, and particularly for many of the inlets that indent the coast, show oxygen supersaturation to be commonly present in the upper 20 m. layer from spring to early f a l l (Data Reports, I.0.U.B.C., 1953-61). The values range normally from 110 to 140% with few exceptions where up to 215% (Data Report No. 5, 1955) and 250% (Gilmartin, 1960) were noted. On the basis of the accumulation of data for these inlets, a general description of the locality of oxygen supersaturation in different water basins can be described. From studies of the general characteristics of British Columbia inlets, Pickard (1961) observed a usual occurrence 63 of two oxygen maxima. The f i r s t , a single surface maximum, was found either near the head or near the mouth of large-runoff inlets, and the second generally appeared below the surface at the mouth or along the length of both large and small-runoff inlets. The latter was usually at a depth of 3 to 8 meters and occasionally as deep as 20 meters, with a location at or below the halocline. Commonly, the surface maximum occurring at the head of the large-runoff inlet dis-played no supersaturation while the other maxima, surface and subsurface, showed supersaturation. The seasonal phenomena and their entailment of large oxygen fluctuations in water initiated the interest for this work. It became the primary concern here to establish the direct relationship between the presence of phytoplankton organisms, their modes of growth and photosynthetic activity, with the extent of oxygen supersaturation that could be reached. It was realized that excellent treatments of prolonged oxygen supersaturation in lakes and inshore waters have already been advanced by Rakestraw (1932), Ricker (1934), Vinberg (1940), Redfield (1948), Hutchinson (1957) and Ramsey (1962). In a l l these treatments, however, supersaturation was considered to be an outcome of biological activity and photosynthesis without any attempt to relate i t to the actual number of organisms present or to their photosynthetic pigment content. The main approach was limited to the study of rates of escape to the 64 atmosphere and the different paths i t follows. It i s the aim, in this second part, to study oxygen supersaturation in sea water in the presence of two common phytoplankton species in terms of (1) their growth and number of organisms produced, (2) the influence of the rate and amount of photosynthetic activity on oxygen production, (3) the i n f l u -ence of their growth and biological activity on the medium, and (4) the dependence of the supersaturation level in water on physical, chemical and biological influences resulting from the presence of these organisms. A review of the background on which these investigations are based is given in order to f a c i l i t a t e the understanding of the reasons which lead to the experimental work, and to the discussion of the results obtained. The production of marine phytoplankton depends on many factors which account for the variation and abundance of cells at different seasons of the year, in various l o c a l i t i e s and at different depths in the ocean. Chemical factors influencing production include the presence of phosphate, nitrate, s i l i c a , oxygen, carbon dioxide and numerous other secondary elements such as iron, sulfur, calcium and manganese as well as salinity and hydrogen ion concentration. Furthermore, organic require-ments of phytoplankton are as important for growth as inorganic substances. Although most of these have been found in the course of investigating the properties of natural substances required to support axenic growth in vitro, nonetheless there 65 is indication for the need of compounds such as diamine tetra acetic acid (EDTA) (Droop, 1957), vitamin B-^ and thiamine (Gran, 1931b). Physical factors involve light, i t s intensity, quality and duration; temperature; viscosity and density of medium. Circulation, particularly upwelling and vertical mixing, are factors that have a great effect upon the supply of nutritive salts (Harvey, 1955, Chap. V). The growth and metabolism of phytoplankton have been studied intensively in laboratory cultures. It is possible under these conditions to give a generalized picture of the growth cycle. Following inoculation of an appropriate medium with the algae and exposure to suitable conditions of light and temperature, growth in cell numbers may begin immediately or after a lag period. Growth is usually exponential in the f i r s t few days, the c e l l numbers per unit volume of medium (n t) at any time (t) being given by the expression (Fogg, 1953) : n t = n Q e r t (12) where n Q is the number of cells per unit volume of medium at zero time, and r the relative growth constant. Following this exponential phase is a period in which the relative growth constant declines continuously and f i n a l l y a stationary phase is reached in which there is no further increase in c e l l numbers. The chemical kinetics underlying the growth sequence in simple algae are presumably similar to those which have been 66 postulated to occur in bacteria (Hinshelwood, 1946). In material which is not actively growing, enzymes may have been denatured and the concentration of essential metabolites may have fallen so that a period of reconstitution i s necessary before exponential growth can begin. In cultures containing a limited volume of medium, exponential growth of algae sooner or later ceases. The factors bringing this about are various, the following being the more important (Fogg, 1953): 1. Exhaustion of nutrient substance or trace element (e.g., Fe) from the medium; 2. Preferential absorption of a particular ion may result in the reaction of the medium becoming inimical to growth; 3. Accumulation of growth-inhibiting products of meta-bolism; and 4. The increasing density of growth may reduce light penetration into the bulk of the medium so that photosynthesis becomes insufficient to maintain exponential growth. The assimilation of organic material in phytoplankton and green plants in general is the direct result of the process of photosynthesis. The energy transfer from light into func-tional chemical energy for the organic build-up is due to green and other accessory pigments present in these plants (Arnon, 1959). These pigments are the green chlorophylls a, b, c, d and e and the various carotenoids, which consist mainly of 67 f -carotene, xanthophyll and xanthopyll epoxide. The carote-noids are found together with chlorophyll in the chromatophores and are present in either an amorphous or crystalline state (Karrer and Jucker, 1950)." The role of carotenoids in photosynthesis has been esta-blished most convincingly in the case of diatoms (Tanada, 1951). The quantum yield has been measured at different wave-lengths and compared with the corresponding estimates of the distribution of light absorption among chlorophylls, fucoxanthin and other carotenoids. The comparison of the quantum yield of photosynthesis with the efficiency curve and those for the distribution of absorption among the pigments in vivo indicates that light absorbed by fucoxanthin is u t i l i z e d in photosynthesis with about the same efficiency as that absorbed by chlorophyll. When carotenoids sometimes occur elsewhere in the proto-plast than in the chromatophores, they are presumably inactive in photosynthesis, and i t may be that within the chromatophore spatial arrangements exclude some carotenoid molecules from participating in active light absorption. Another possibility is that certain types of the carotenoids are able to absorb light for use in photosynthesis but that others are quite inactive in this way (Fogg, 1953). Further information about the part played by accessory pigments in photosynthesis comes from studies of fluorescence. The yield of fluorescence, i.e., the ratio of the energy emitted to that absorbed, gives information regarding the fate of the excitation energy. The yield of the chlorophyll fluor-escence is the same whether i t is excited by red light (6000R), absorbed exclusively by chlorophyll, or by blue-green light (4700X), three-quarters of which is absorbed by carotenoids (Dutton et a l . , 1943). This evidence of transference of energy from accessory pigments to chlorophyll suggests that light absorbed by the former pigments is not u t i l i z e d directly in the photochemical reaction. It appears, in fact, that a l l light energy used in photosynthesis must pass through chloro-phyll-a. Thus, in Chlorella only chlorophyll-a fluoresces and it s fluorescence is excited by light absorbed by chlorophyll-b (Duysens, 1951). Since chlorophyll-a is the principal and only photosynthetic pigment common to a l l the algal classes this provides evidence that the photochemical reaction i s of the same nature in a l l algae. If chlorophyll-a is the essential catalyst for photo-synthesis, i t might be reasonable to assume that the gross rate of photosynthesis taking place in sea water illuminated by light of a given intensity is a function of chlorophyll-a content of sea water. There is evidence, however, that caro-tenoids also influence this relationship. Yentsch and Ryther (1957) showed that a linear relationship holds between the ratio of net to gross photosynthesis and the ratio of chloro-phyll-a to the total carotenoids in a water sample. 69 A diurnal fluctuation in the rate of photosynthesis by a plant in i t s natural surroundings i s expected, owing to fluctuation in the light intensity to which i t is exposed. In addition there seems to be also a diurnal or similar varia-tion in the intrinsic photosynthetic activity or'Vitality" of the plant i t s e l f (Strickland, 1960). Rabinovitch (1951), Doty and Oguri (1957) and Ryther (1956) have provided informa-tion on the existence of diurnal, geographical and annual influences on rates of photosynthesis. The phenomenon was given the name "Periodicity" for i t s consistently maintained pattern under various conditions. By contrast, the simple unicellular algae such as found in marine phytoplankton can be grown in culture for many days or weeks with no notable daily periodicity in growth rate after an i n i t i a l adjustment period (Strickland, 1960). The process of growth of the organisms themselves could cause certain changes in the water or medium in which they l i v e . These changes may act as inhibitive factors for further growth as well as for photosynthesis. For example, carbon dioxide concentration (and pH level) might prove a limiting factor in the supply of reductant material (Arnon, 1959). Tailing (1960) indicates that inadequate carbon dioxide supply and inadequate buffering do not affect photosynthesis for short exposures, but might do so after long exposures when culture growth is increasing to high densities. Furthermore, the 70 release of extracellular products from li v i n g , dead or decom-posing plankton could inhibit photosynthesis or the growth rate by the different reactions they may undergo in either culture media or at sea. Saunders (1957) sums up these effects in four major points: 1. The extracellular products may supply compounds ut i l i z e d as energy sources or which contain the basic elements essential to the building of protoplasm. 2. They may supply accessory growth factors which are required for or merely stimulate the growth of organisms. 3. They may be in the form of toxic substances which inhibit a c t i v i t i e s or cause the death of many aquatic organisms. 4. They may form organic complexes with trace metals, the effect of which may be beneficial or detrimental, depending upon circumstances. Matsudaira (1950, 1951, 1952a) studied the effect of the catalytic activity of sea water on the decomposition of hydrogen peroxide in the presence of organic matter similar to that extracted from f i l t e r e d culture water of different age. He found a substantial decrease in the decomposition of hydrogen peroxide in both cases, the catalytic activity decreasing rapidly with higher concentrations of either organic matter or extracellular product. This is equivalent to catalytic inhibition. The v e r s a t i l i t y of this method in testing for some of the medium influences on growth 71 recommends i t for further use and study in this connection. The variation of the rate of desupersaturation at sea or in culture water is considered to be influenced by two major factors: the f i r s t i s a physical-chemical one, which is related to ion content, surface tension, temperature, agitation, surface area of water as related to the "effective" exposed surface, depth of supersaturation column and other factors that contribute to nucleation and spontaneous formation of gas bubbles. The second factor is mainly an outcome of the presence of phytoplankton and other organisms in the super-saturated water. These organisms may act in various ways to reduce the supersaturation level. Chiefly, they may consume oxygen by respiration and oxidation, provide a "wall" effect for nucleation, and even during photosynthesis their average closeness to each other can provide the steps for bubble formation and rapid escape. By excluding the biological influence on oxygen in water after a rapid and vigorous interval of photosynthetic activity, the mechanisms of desupersaturation can be classified relative to their exchange with the atmosphere in the following manner. 1. For an undisturbed water surface of a quiescent water body, exchange with the atmosphere is governed by an uppermost thin layer (Hutchinson, 1957) which i s at equilibrium with the atmosphere above i t , any escape or entrance of gas such as oxygen from and to the water body is governed by Fick 1s diffusion 72 equation from that layer. T T - = D T T " < 1 3 > Where Z indicates depth, t time and D the molecular diffusion coefficient. D at 20°C. i s of the order of 2.0x10""' cm2 sec! 1 and i t s rate of change with temperature is about 3% deg 1 C. (Millington, 1955). 2. For a well mixed body of water where the concentration of dissolved oxygen i s uniform through the water at any time, the rate of movement from and to the liquid of oxygen across the surface depends on the difference between the actual concen-tration and the saturation value. The rate of entrance and escape is given by Bohr's equation (Bohr, 1899). d °2 -dT" = a * <P"Pt> <14> Where a is the surface area; cK , is the entrance coefficient; P, partial pressure of gas in the atmosphere; and p t = where t designates time and s, saturation value at the tempera-s' )°2^t ture and ion content. A modification of this equation was used by Dorsey (1940, page 564) and Redfield (1948) to interpret the escape rate of an escape coefficient ^ which contains — - . j% then is oxygen from supersaturated water. This was done by introducing L i >2 given by the following expression: fi - - ' I l n f l - *> °2}t j ( 1 5) L ZJo p where "Z is average depth of water column; t, time; L A O o l 73 difference between supersaturation and saturation at time o and t. The above treatments are advanced essentially for dis-solved oxygen molecules assuming that only individual molecules take part in the exchange, where the rate of exchange is believed to follow an exponential time dependence (Findlay and King, 1913). A more important path of oxygen escape in supersaturated waters, especially pertinent to the sea and other natural water bodies, may occur through formation of gas bubbles. These bubbles rise to the surface and escape to the atmosphere in a relatively short time compared to that taken for molecular exchange. It is believed that microbubbles which act as nuclei for formation of larger gas bubbles may result from the bio-logical activities of photosynthesis and respiration, although there is no experimental proof of this mechanism (Ramsey, 1962). Theoretical treatment for bubble formation has been dealt with by Harvey et a l . (1944). They attributed the main contri-butions to bubble formation to difference between total pressure on the bubble and oxygen tension (P-pfc) as defined for Bohr's equation (14). However, P is taken here as the sum of atmos-pheric and hydrostatic pressure over the bubble. Other contri-butions include sites for growth of nuclei, diffusion and surface tension of the water. Miyake (1951), Hutchinson (1957) and Ramsey (1962) followed by treatments of conditions 74 of allowable bubble formation in the sea. Their calculated results were based on ion content, temperature, and total pressure as factors determining the capacity of water for saturation. Earlier, Dean (1944) considered bubble formation and cavitation as a result of movement in liquids. These occur in vortices produced by turbulent motion of the liquids. The vortices could expand bubble nuclei on foreign bodies and dust particles, i f they were present, and hence largely con-tribute to bubble formation and escape. The st a b i l i t y of the bubble in the liquid after i t s formation depends on pressure and the presence of other gases in solution. The capability of one gas to act as a "seed" for bubble formation renders the bubble a mixture of a l l the gases found in the liquid. Wyman (1952) related the stability of the bubble to concentration of the different gases, the hydrostatic pressure and the relative diffusion coefficients of each gas. He showed further that the size and the relative concentrations of gases in the bubble is subject to a continuous change which accompanies i t s ascent to the surface. In order to establish the direct relationship of oxygen supersaturation in sea water with photosynthetic activity of phytoplankton and i t s variation with growth characteristics, experiments were set up to study unialgal cultures of two commonly occurring species of marine phytoplankton in the laboratory. Laboratory experiments were preferred to f i e l d investigations because important parameters which influence growth such as light intensity, temperature and nutrient con-centration, could be controlled. This was f e l t to be necessary due to the restriction and limitations which may be imposed on a study of this nature in the sea. The continuous variation of growth-controlling parameters in natural water bodies which include temperature, light, circulation as well as predation by animals are but few of the factors that could interfere in assessment of any obtainable result. Furthermore, the time of investigation is restricted to the time and location of the appearance of a phytoplankton bloom, which is usually not readily detectable without some search. For the purposes of discussing the results, the proposed experiment is divided into three parts: The f i r s t includes continuous observation of oxygen and pH in culture water relative to number of organisms, concentration of photosyn-thetic pigment and the influence of changes in the medium on growth. The second part deals with fluctuation and extent of oxygen supersaturation under various intervals of light and dark periods. The third takes into consideration the behaviour of organism-free oxygen supersaturated water and the effects of salinity, surface tension, s t i r r i n g and geo-metry of the water column on desupersaturation rates. 76 EXPERIMENTAL I. Methods and Apparatus 1. Oxygen Determination (i) Winkler method: (After Strickland and Parsons, 1960). Concentration of oxygen is reported as ml. 0 2 (at STP)/1. of water. The experimental error is assumed at - 0.5%. ( i i ) Microgasometric technique: (After Scholander et a l . , 1955). Both oxygen and nitrogen are reported as ml. 0^ or H (at STP)/1. of water. The estimated error for this method is * 3-4%. ( i i i ) Electrode system technique: (After Carrit and Kanwisher, 1959). Determination of dissolved oxygen was made with a "Jarrel-Ash Co. Model No. 26-601 Dissolved Oxygen Analyzer" consisting of a silver-silver oxide reference elec-trode and platinum indicator electrode. The principle in-volved the measurement of oxygen diffusion through 0.0005" teflon membrane wrapped around the electrodes with a weak solution of potassium hydroxide placed in between the membrane and the electrodes. Since diffusion to the electrode is a measure of the concentration of oxygen, the st i r r i n g of the water was an essential requirement. The concentration of oxygen in water is reported as percentage or as ml. 0 2 (at STP)/1. of water. Calibration with the Winkler method indicates 77 . an error of * 37o. 2. Salinity Determination: (After Strickland and Parsons, 1960). A 10 ml. sea water sample was titrated with a standard-ized, approximately 0.2 N silver nitrate solution using potas-sium chromate as an end point indicator. 3. pH Measurement and Total Alkalinity: (After Strickland and Parsons, 1960). The pH and total alkalinity were determined using a "Beckman Model GS pH Meter". pH was determined on the sample of sea water, and total alkalinity on 100 ml. sample of sea water before and after the addition of an exact quantity of 0.0100 N hydrochloric acid. 4. Light Intensity and Radiation Energy: The light intensity used to illuminate the cultures was determined by a "Photovolt Electronic Photometer - Model 501-M", which gave the intensity of the visi b l e region in foot-candles. This was converted into radiation energy using the factor 1 ft.-cdle. = 2.3x10"5 langly/min. (langly/min. - gram-cal/cm2/min.) (Strickland, 1958). 5. Cell Count: From a well shaken culture, 1 ml. of medium was taken with a pipette and placed in a "No. 800 Sedgewick-Rafter Counting Chamber". The dimensions of the chamber were 5.0x2.0x0.1 cm. The liquid was covered with a thin cover s l i p , and l e f t for 15 minutes or u n t i l the orga-nisms settled completely to the bottom. The chamber was placed under the microscope, and with the aid of a calibrated 78 eyepiece, five counts were performed over different regions of the chamber. The counts were averaged and multiplied by the calibration factor. The count is reported as No. of organisms/1, of water. 6. Pigment Concentrations: Chlorophyll-a, b, and c and Carotenoids (After Richards and Thompson, 1952, and Creitz and Richards, 1955). The method consisted of f i l t e r i n g a quantity of a well shaken, homogeneously distributed culture medium through a "millipore HA, 0.45 pore size" f i l t e r . The f i l t e r was placed in 5 ml. of 90% acetone for 20 hours. The liquid was then transferred to a 10 ml. centrifuge-tube and centrifuged for 15 min. Three or four mis. were trans-ferred afterwards to a 1 cm. path-length c e l l and immediately read in a "Beckman D.U. Spectrophotometer" against a c e l l containing acetone at 665 raj*, 645 mjfi, 630 myi, 510 mp. and 480 mji. Concentrations of chlorophyll-a and b are reported in mg./l. and chlorophyll-c and carotenoids in mspu/1. 7. Rates of Photosynthesis and Respiration: (After Strickland and Parsons,1960). Part of a well shaken culture sample was transferred into 2 - 150 ml. light B.O.D. bottles and 2 - 150 ml. dark B.O.D. bottles. Oxygen concentration was determined on the other part of the i n i t i a l sample and on the B.O.D. bottles after one to three hours from incubation under the same light intensity and temperature. Rates of photosynthesis and respiration are reported as mg. C/l./hr. 79 8. Catalytic Activity of Water on K^Oo: (After Matsu-daira, 1950). A thirty per cent hydrogen peroxide solution was diluted to 1/100 with double d i s t i l l e d water. This was mixed well and kept in the dark at 12°C. A 5.0 ml. of the stock solution was transferred to 50.0 ml. of sample in a 125 ml. Ehrlenmeyer flask and to 50.0 ml. of double d i s t i l l e d water used as a standard. The flasks were incubated for 24 hours at 12°C. after which they were removed and 10 ml. of each was transferred to a flask containing 10 ml. of d i s t i l l e d water, 1 ml. of (1:3) sulphuric acid, 0.5 g. KI and 1 drop of 10% solution of ammonium molybdate. The flasks were l e f t for 3 minutes and then the liberated iodine was titrated with a 0.009 N sodium thiosulphate solution. A small modifi-cation of Matsudaira's method was used by taking the catalytic decomposition as the difference between the d i s t i l l e d water standard and the sample instead of the concentration in the sample before and after incubation. Catalytic activity was reported as Kcat. hr."*-, the reaction rate constant of a f i r s t order decomposition rate. 9. Surface Area of Diatomaceous Earth: (After Brunauer, Emmett and Teller, 1938). Surface area was determined with spectroscopic nitrogen at various pressures and was found to be 1.03 m2. gm-1. 80 10. Surface Tension Measurements - Capillary Rise  Method: (After Daniels, Mathews et a l . , 1956, p. 52). The capillary radius was calibrated with double d i s t i l l e d water using 71.80 dynes cm"1, at 25.0°C. A l l measurements were done in a thermostated bath with a Fisher apparatus No. 14-817 (1961). 11. Enriched Sea Water for Culture Growth: Sea water ranging in salinity between 25.07<> to 33%©was fi l t e r e d through an HA millipore f i l t e r of pore size 0.45 yu. To the f i l t r a t e , the following ingredients were added in relative concentrations to one l i t e r of water (Sweeney, 1961); Sodium Nitrate 0.25 gm. Potassium Monophosphate 0.02 gm. Sodium Silicate 0.10 gm. Trade Element Mixtures - Iron 0.5 mg. Zinc 0.3 m^g. Manganese, Molybdenum, Boron, Cobalt, Copper 0.1 mg. Thiamine 0.5 gm. II. Reagents 1. Phytoplankton Species (i) Nitzschia closterium: An agar slope of this diatom strain was received in February, 1961 from Scripps Institution of Oceanography. It was subcultured immediately into s t e r i -lized 125 ml. Ehrlenmeyer flasks containing 50 ml. of enriched sea water under aseptic conditions. 81 ( i i ) Chlorella Strain "A": A broth strain was received in March^ 1961 from Woods Hole Oceanographic Institution and subcultured as above. (This strain is toxic to clam and oyster larvae: U. S. Fish and Wildlife Survey Vol. 58, 1958.) 2. Diatomaceous Earth: " J . T. Baker" Reagent grade. 3. Chemicals: A l l chemicals used were reagent grades except "spectranalyzed" acetone. III. Procedure A (i) Observation and measurement of oxygen concentration and pH in cultures of Nitzschia and Chlorella. Two glass troughs 26x16x20 cm. were cleaned thoroughly with n i t r i c acid, rinsed with d i s t i l l e d water, and then wiped with methyl alcohol and dried. Each trough was f i l l e d with 3 l i t e r s of steri l i z e d and enriched sea water of salinity 27.017o. One trough was inoculated with Nitzschia closterium and the other with Chlorella strain "A". The troughs were placed on a shelf of an especially built unit under fluorescent "cool white" light source. Measured radiation energy reaching the water surface was of the order of 9.2x10 ly./min. By means of an automatic control the light period was adjusted to 16 hours continuous illumination followed by 8 hours of darkness. Measurements of oxygen, nitrogen and pH were made every few days, exactly 6 hours after the beginning of the light period. Oxygen and nitrogen concentrations were deter-82 mined by the gasometric technique of Scholander et a l . (1955). A ( i i ) Estimation of pigment concentration, number of organisms, respiration and photosynthetic rates, pH and total alkalinity in a culture of Nitzschia closterium. Nine plastic (methyl methacrylate) troughs, 26 cm. in diameter and 9 cm. high, were cleaned as before and each f i l l e d with 1 l i t e r of ste r i l i z e d and enriched sea water of salinity 27.017o. Each trough was inoculated with a 2 ml. suspension from a well mixed subculture strain of Nitzschia. These were incubated at 12^1°C. under the same light and dark condi- , tions of A ( i ) . Oxygen, nitrogen and pH were determined every few days, 6 hours after the start of light period. Consecu-tively, every few days, one trough was taken out and i t s contents transferred to a 2-liter flask and shaken well to (a) give a homogeneous distribution of organisms in the medium and (b) to bring the oxygen concentration near the saturation level. Part of the sample was used to determine photosynthesis and respiration rates, number of organisms, total alkalinity and salinity. The other part was f i l t e r e d through a millipore f i l t e r and the f i l t e r was used to determine pigment content of the sample. A ( i i i ) Determination of catalytic activity of culture water by rate of decomposition of hydrogen peroxide. From two troughs as in A ( i ) , one inoculated with Nitzschia and the other with Chlorella and each containing 4 l i t e r s of 83 enriched sea water, salinity 25.5%, two 50.0 ml. samples were taken from each concurrently with oxygen, nitrogen and pH determination. The samples were f i l t e r e d into extra-clean 125 ml. Ehrlenmeyer flasks, 5.0 ml. of hydrogen peroxide stock solution were added to each, and to a d i s t i l l e d water standard; the samples and d i s t i l l e d water standard were incubated at 12^1°C. for 24 hours. At the end of the incuba-tion period, the samples and standard were titrated with thiosulphate solution after addition of potassium iodide and sulphuric acid and the rate of decomposition computed from the difference in concentration between standard and sample. Controls for different pH levels were also made on a sample of the original medium before inoculation, and the variation due to pH was used to correct for the results. The results are reported as Kcat. hr.""'-. B (i) Study of light and dark period effects on the oxygen level in a heavy culture of Nitzschia closterium. A plastic (methyl methacrylate) tank 44x24x60 cm. was cleaned and h a l f - f i l l e d with 32 l i t e r s of enriched and f i l t e r e d sea water, salinity 30.20%. The water was inoculated with Nitzschia closterium and placed in the incubator under the same light condition as before. After 22, 45, and 90 days, oxygen concentration was measured at different depths by means of glass tubes set at different levels in the tank. The samples were extracted by siphoning the water through a glass tube into a B.O.D. bottle when the Winkler method was used, or the electrode of the oxygen analyzer was employed by moving i t vertically in slow motion at the desired depth. Measurements were taken at the beginning of either a light or dark period and at different time intervals afterwards for a total period of not less than 5 hours. C (i) Measurement of desupersaturation rate constant in d i s t i l l e d and sea water under quiescent conditions. In a similar tank as B (i) was placed 32 l i t e r s of fi l t e r e d sea water, salinity 32.00%„. Oxygen gas was bubbled through a diffuser stone in the water at 0.5 l i t e r min. 1 for 4 minutes or u n t i l oxygen supersaturation exceeded 150%. The tank was placed in the incubator at temperature of 12^1°C. Oxygen concentration was determined at various depths by means of the Winkler method at intervals of 7 to 8 hours. The procedure was repeated for smaller quantities of water in the tank, and also for d i s t i l l e d water. Relative humidity ranged from 84% to 95% in the thermostated room. The desuper-saturation rate constant was calculated from the integrated f i r s t order rate equation and reported as Kdes. hr. 1 C ( i i ) Measurement of desupersaturation rate constant for waters of different s a l i n i t i e s under stirred conditions. One l i t e r of water of salinity 32.80%«>was placed in a 2-liter glass beaker containing a teflon-covered magnet bar. The water was supersaturated by bubbling oxygen gas through i t 85 as in C (i) to a level not exceeding 150%. The oxygen analyzer electrode was placed in the middle of the beaker 2 cm. below the surface, and the magnetic stir r e r operated at about 800 r.p.m. Oxygen concentration was read on a recorder chart at different intervals of time. The rate constant, Kdes. hr."**' was computed as before. The same procedure was repeated for water ranging in salinity between 0.00 and 32.80%*. C ( i i i ) Measurement of desupersaturation rate constant in presence of various quantities of diatomaceous earth or different concentrations of surface-active agent (heptanoic acid) under stirred conditions. To the same tank used in C (i) containing 32 l i t e r s of sea water, salinity 27.707o<»and two magnet bars operated by two stirrers at 800 r.p.m. was added a weighed amount of diatomaceous earth. The water was supersaturated by bubbling oxygen gas through i t , and the rate of desupersaturation was measured at different time intervals totalling not less than 7 hours. The same procedure was repeated for 32 l i t e r s of sea water salinity 33.01%©with addition of different quantities of heptanoic acid. The surface tension was measured after each such addition in a thermostated bath at 25.0°C. The desuper-saturation rate constant Kdes. hr."*" is reported for the total surface area of diatomaceous earth, and for variation in surface tension. 86 RESULTS AND DISCUSSION A (i) Variation of oxygen concentration and pH in cultures of Nitzschia and Chlorella. Oxygen concentration and pH were determined concurrently at the fixed time of 6 hours after the beginning of the light period in order to relate photosynthetic activity to changes in these quantities in a sufficient and given time. It is well known, however, that better growth occurs under aeration and shaking of cultures. This is presumably due to the accele-ration of carbon dioxide diffusion into the water which buffers the medium and lowers the pH. Nonetheless, an aerated or shaken culture could neither represent a meta-stable system of a dissolved gas where concentration i s a function of influx and escape, nor define the limitation of pH as an inhibitive factor. It i s assumed, therefore, in maintaining quiescent conditions for these cultures, that at least the influence of outside sources i s minimized. Oxygen, nitrogen and pH values for the two cultures are given in Table XIX and shown graphically in Figure XX for Nitzschia and Figure XXI for Chlorella. It could be seen from these presentations that oxygen and pH follow the same pattern of change, indicating that both are the result of photosynthetic activity and carbon assimilation. The values 87 TABLE XIX VARIATION OF 0 2, N 2 AND pH IN CULTURES OF NITZSCHIA AND CHLORELLA Exp. A ( i ) , Temp. 12*1°C, S%»= 27.01, 6 hours light Oxygen Saturation Level = 6.17 ml./I. I n i t i a l Concentration: Nitzschia 10 6 org./l. Chlorella 4xl0 6 org./l. Time Days 0 2 ml/1 N 2 ml/1 pH 0 2 ml/1 N 2 ml/1 pH 1 5.75 10.75 8.19 6.25 11.00 8.05 3 6.75 11.13 8.23 6.50 10.88 8.25 5 .9.00 11.75 8.60 6.88 11.50 8.32 7 11.00 11.13 9.10 7.87 11.00 8.50 9 12.63 11.38 9.55 8.38 11.50 8.75 10 11.38 11.00 9.63 8.68 11.25 8.77 12 10.88 11.00 9.45 9.63 11.00 8.85 14 11.00 11.00 9.65 9.50 11.13 9.22 15 11.75 11.00 9.70 9.13 11.00 9.25 19 11.25 11.00 9.62 9.00 11.00 9.48 21 11.75 11.00 9.87 9.00 11.13 9.63 23 13.00 11.37 9.65 10.50 11.37 9.49 26 11.00 11.38 10.00 9.00 11.88 9.35 28 10.25 11.38 9.85 9.25 11.63 9.50 30 11.75 11.63 10.05 10.25 11.63 9.40 33 10.38 11.50 9.70 10.00 11.63 9.50 35 10.25 11.63 9.60 10.50 11.63 9.40 37 9.25 11.63 9.55 9.50 11.63 9.50 40 9.75 11.25 9.60 10.50 11.63 9.55 42 9.38 11.63 9.45 10.38 11.63 9.55 44 8.50 11.63 9.45 9.75 11.63 9.60 Oxygen and nitrogen measured by micro-gasometric technique increase rapidly in the f i r s t few days, after which they fluc-tuate about a maximum level for a period of time reaching over 30 days. Then both values start showing signs of decline in Nitzschia, but not in Chlorella. Nitzschia was by far the more active of the two species, producing oxygen concentrations in water in excess of 200% saturation and pH up to 9.00. 88 In comparison, the values of oxygen and pH in the Chlorella culture were about 30% and 15% less, respectively. Estimation and measurement of nitrogen in a photosynthetic medium serves many purposes. On the whole, nitrogen in the sea does not participate in any reaction that could noticeably alter i t s concentration. Furthermore, i t is quite distinct from either carbon dioxide or oxygen, because of i t s inab i l i t y to enter into any photosynthetic process. This rather inert nature of nitrogen was conveniently put to use by oceanographers as an indicator of the extent of gas exchange between atmosphere and hydrosphere. The observed concentrations of nitrogen in both cultures, Figures XX and XXI, are seen to assume a rela-tively constant level with time. However, on the whole, these concentrations were about 1.0 to 1.5 ml/1 less than the satura-tion values for that particular salinity and temperature (Hamm et a l . , 1941). The rapid production of oxygen under these conditions i s expected to displace and "flush" nitrogen out of the medium in the course of i t s escape to the atmosphere. It was assumed by Redfield (1948) that in l o c a l i t i e s of oxygen supersaturation in the sea, two-fifths of the total oxygen exchange with the atmosphere is due to oxygen produced by photosynthesis. Hence, use of nitrogen as an indicator for such exchange could not be totally applicable in these regions, at least not during the seasonal occurrence of oxygen super-saturation . Similar results to those found here had been obtained at sea. In their study of the dynamics of a phytoplankton bloom, Ryther et a l . (1958) reported that both oxygen and pH followed a similar course of change during periods of photosynthetic activity. pH varied between 8.5 early in the morning to 9.5 in the afternoon. They assumed that at\its maximum no free carbon dioxide was available and any further photosynthesis was dependent upon bicarbonate and carbonate ions. The water was approximately 90% saturated with oxygen at daybreak, after which i t increased steadily t i l l i t reached a maximum in the early afternoon. One such maximum was reported to have reached 270%. A ( i i ) Relation of growth and photosynthetic activity to oxygen supersaturation in Nitzschia. a. Number of organisms Frequent determinations of oxygen, nitrogen and pH were carried out on nine separate cultures of Nitzschia incubated at a fixed time interval of illumination in order to account for any deviation which might take place in any one culture apart from the rest. Table XX shows that oxygen and pH were similar at any day's measurement for a l l cultures within a few per cent of the average values. This agreement allows for treatment of a l l these separate cultures, being under the same conditions, as one system; further, i t permits the termina-tion and study of one certain culture at a time without 90 TABLE XX SERIES CULTURE OF NITZSCHIA Temp. 12*1°C, S%e= 27.01, 6 hours light I n i t i a l Concentration: 58x10 organisms/1. B l B 2 Time Days 0 2 ml/l N 2 ml/l pH 0 2 ml/l N 2 ml/l pH 1 7.13 12.00 8.55 7.13 12.00 8.50 2 7.50 12.50 8.55 7.13 12.00 8.50 4 8.50 11.63 9.00 6 9.38 11.88 10.10 B3 B4 Time Days 0 2 ml/l N 2 ml/l pH 0 2 ml/l N 2 ml/l pH 1 7.00 11.88 8.50 7.13 12.00 8.63 2 7.75 12.37 8.50 7.25 12.13 8.50 4 9.28 11.88 9.10 8.75 11.88 9.23 6 8.33 11.50 10.38 8.50 12.00 10.50 8 8.75 11.63 10.50 8.33 11.63 10.40 10 8.00 11.50 10.50 13 7.50 11.87 10.50 B r B, 5 6 Time Days 0 2 ml/l N 2 ml/l PH 0 2 ml/l N 2 ml/l pH 1 6.88 12.00 8.55 7.13 12.00 8.65 2 7.38 12.13 8.50 7.25 12.00 8.42 4 8.00 12.00 8.85 7.63 12.00 8.68 6 7.38 11.75 9.22 7.25 11.63 8.82 8 8.25 11.63 9.40 7.00 11.63 8.70 10 8.13 11.63 9.73 8.00 11.88 9.00 13 7.37 11.75 9.80 7.63 11.63 9.00 15 6.88 11.75 ' 9.75 6.88 12.00 9.02 17 7.25 11.63 9.15 (continued) 91 TABLE XX (Continued) B„ ' B 8 Time Days 0 2 ml/1 N 2 ml/1 pH 0 2 ml/1 N 2 ml/1 pH 1 7.25 11.88 8.70 7.13 11.50 8.75 2 7.63 12.00 8.70 7.50 11.63 8.78 4 9.00 11.88 9.47 8.88 11.75 9.47 6 8.37 11.88 9.60 9.00 11.88 9.85 8 8.38 11.63 9.45 8.63 11.75 10.20 10 8.50 12.13 9.40 7.75 11.88 10.27 13 7.50 11.88 9.70 7.25 11.88 10.20 15 7.00 11.88 9.85 7.50 11.88 9.95 17 7.50 11.63 10.45 7.50 11.75 10.10 20 6.38 11.50 10.35 7.75 11.63 10.35 22 7.00 11.63 10.25 B 9 Time Days 0 2 ml/1 N 2 ml/1 PH 1 7.13 11.25 8.58 2 7.50 11.63 8.48 4 7.63 11.50 8.88 6 7.00 12.38 9.10 8 7.63 11.75 9.20 10 8.00 11.88 9.40 13 7.75 11.88 9.35 15 7.38 12.00 9.22 17 7.50 11.63 9.30 20 7.13 11.63 9.25 22 6.20 11.38 9.25 Oxygen and nitrogen measured by micro-gasometric technique disturbing the system as a whole. A l l essential quantities that were anticipated to influence or control oxygen super-saturation in these cultures are given in Table XXI. These include number of organisms, photosynthetic pigment content, rates of photosynthesis and respiration and total alkalinity. TABLE XXI QUANTITIES INFLUENCING OXYGEN PRODUCTION IN A NITZSCHIA CULTURE Sample No. B l B2 B3 B4 B5 B6 B7 B8 B9 Age after incubation 2 6 8 13 15 17 20 22 23 Chlorophyll-a mg/1 34 184 112 121 76 73 130 80 163 Chlorophyll-c mspu/1 28 86 130 117 181 93 209 134 174 Carotenoids non-astacin mspu/1 15 62 61 67 54 57 100 74 116 Number of organisms x l 0 " 6 / l 171 622 524 446 915 865 1063 809 1076 Total al k a l i n i t y 2.13 2.52 2.34 3.07 2.83 2.69 2.96 2.96 3.07 Rate of photosynthesis mg C/l/hr 0.290 0.912 0.358 0.031 0.156 0.072 0.201 0.040 0.192 Rate of respiration mg C/l/hr 0.059 0.209 0.247 0.237 0.402 0.178 0.230 0.316 0.284 (continued) TABLE XXI (Continued) Sample No. B l B2 B3 B4 B5 B6 % Saturation of oxygen 121 152 142 122 111 117 Net Gross photosynthesis 0.83 0.84 0.64 0.14 0.32 0.33 ™> ^ (Chlorophyll-a - Carotenoids)19.6 98.1 51.0 54.4 21.8 15.0 B7 103 0.66 1.30 30.1 B8 113 0.13 1.08 5.9 B9 100 0.44 1.40 46.5 94 The total number of organisms reached in each culture had been expected to be an indication of the growth rate and, hence, to some extent, related to the degree of biosynthesis. However, Eigures XXII and XXIII show that the number of organisms is not directly proportional to photosynthetic pigment concen-tration nor to rates of photosynthesis and respiration. On this evidence, c e l l counts may then be dismissed as indices to the relative degree of biosynthesis in a given sample or to growth rate after the stationary phase of growth is reached. A similar conclusion to this effect was reached by Graham (1943) on the basis of the usually extreme difference in skeleton size or skeleton weight found among the organisms. It is generally believed, however, that chlorophyll and auxi-l i a r y pigment assays should afford a more reliable measure of the potential organic synthesis in any given sample or region. There are many indications that chlorophyll concentration as well as carotenoids vary diurnally for a certain number of organisms, depending on the light intensity (Yentsch and Ryther, 1957; Ryther et a l . , 1958; Yentsch and Scagel, 1958) and other growth factors. This is independent of new and dead cells variation, but more a change of the stationary phase. Though the organisms may not be actively growing they are not necessarily either dying or perishing in great numbers. 95 b. Photosynthetic pigment concentration. It was logical to relate oxygen supersaturation to the magnitude and duration of net photosynthesis in a medium, where net photosynthesis was the excess oxygen produced in presence of respiration and oxidation. The rate of net photo-synthesis in turn should be dependent on the concentration or the activity of photosynthetic pigment under constant conditions of light temperature and sufficient nutrient concen-trations. Since chlorophyll-a was considered to be the chief agent for conversion of light energy into chemical energy, the primary step was an attempt to compare concentration of chlorophyll-a to rate of net photosynthesis. The relation was non-linear pointing to the possibility of yet another factor which might enter and influence oxygen production. By correlating carotenoid concentration, which i s assumed to be a polyfunctional pigment with chlorophyll-a, a linear rela-tion was obtained by plotting the difference of the two pigment concentrations (chlorophyll-a - non-astacin carotenoid) versus rate of net photosynthesis, Figure XXIV. An approximately similar result was obtained by Yentsch and Ryther (1957) and Ryther et a l . (1958) by employing the following relationship chlorophyll-a gross photosynthesis R carotenoids = 1 net photosynthesis x 2 ^ ' Comparison of Figure XXIV to the relationship applied by Yentsch and Ryther, Figure XXV shows that both relationships behave approximately in a similar manner. Although i t appears that the application of difference in pigment concentrations f i t s the results with less point scatter than the ratio of pigment concentrations, nevertheless, the influence of caro-tenoids on net photosynthesis cannot be neglected. Carotenoid pigments have been assumed to play multiple and various roles in the c e l l , chiefly among these is the transfer of light energy to chlorophyll in diatoms at about 4700A* (Tanada, 1951; Dutton et a l . , 1943) when the pigments are in the chromatophore. Otherwise certain of these caro-tenoids are inactive either because of some chemical change in their composition or because of their occurrence outside the chromatophore (Fogg, 1953). In this latter case they most probably act as light scavengers and shields to light penetra-tion to the chromatophores. The carotenoids also contain a large number of double bonds which form conjugated systems giving rise to the possibility of a large number of cis-trans isomers, many of which have been isolated. Retrovsky (quoted by Goodwin, 1953) suggested that fucoxanthin together with violaxanthin and zeaxanthin play a part in redox system in algae. This latter suggestion is shown to apply here by the comparison of the rate of respiration to concentration of carotenoids, Figure XXVI. Further c l a r i f i c a t i o n of the complex roles of carotenoids i s made by Tanada (1951) who showed that only light absorbed by fucoxanthin, which i s one of five 97 xanthophyll pigments in diatoms is u t i l i z e d by photosynthesis in the same efficiency as light absorbed by chlorophyll. In the case of the remaining carotenoids, however, light absorbed by them appears not to be used in the photosynthetic process. Pace (1941) found that in Nitzschia culture the average ratio of chlorophyll to carotenoids was 3.77 and that of xanthophylls to carotene to be 8.35. Assuming that fucoxanthin is one-third of the xanthophylls (Strickland, 1960), then one-fourth of the carotenoids could participate actively in photosynthesis. The linear relation of difference in chlorophyll-a and caro-tenoids to rate of photosynthesis indicates the possibility that the rest of the carotenoids may act as a light shield for the chlorophyll. Light intensity, furthermore, has a direct influence on photosynthesis not only by the light quanta offered to the pigments but also by the manner i t conditions both chlorophyll and carotenoids concentrations in the c e l l . There i s evidence that light intensity and duration directly influence pigment build-up in the sea. Diurnal investigation of pigment concen-tration relative to light intensity by Ryther et. ajL. (1958) in Senix Creek showed chlorophyll concentration to have increased gradually during the day u n t i l i t reached a peak at sunset, after which i t decreased rapidly throughout the night u n t i l daybreak. Also, Emerson (1929), Fleisher (1936), Sargent (1940) and Myers (1946) reported chlorophyll concentration to 98 have increased with light intensities from 2 to 12-fold range in Chlorella c e l l s . Similarly plant carotenoid pigments increased during the day and decreased at night, (Ryther et a l . , 1958). The ratio of chlorophyll-a:carotenoid decreased through-out the daylight periods from a maximum observed at sunrise. The change in chlorophyll is usually much larger than in carotenoids. Spoehr and Milner (1949) observed a 10-fold-change in chlorophyll to the magnitude of change in carotenoids for Chlorella at high cellular yield. No correlation of chlorophyll-c to any apparent activity seemed feasible under the experimental conditions, although theoretically chlorophyll-c i s supposed to transmit light energy to chlorophyll-a acting as an accessory pigment. Yentsch and Seagel (1958) assume on the basis of Granick's work (1949) that chlorophyll-c is a percursor pigment of chlorophyll-a. The actual interpretation of chlorophyll-c is l e f t , however, for further studies and investigation. Because of the constant conditions of light intensity, duration and temperature employed in this experiment, the rate of net photosynthesis and hence the difference in pigment con-centration was expected to be proportional to the amount of oxygen saturation found in the culture medium at the time of rate determination. Figure XXVII shows per cent saturation to be independent of rate of respiration but in a manner related to rate of photosynthesis. The pattern of the increase in both 99 quantities i s apparent but not to the expected degree. This, however, is not surprising when i t is remembered that the condi-tions under which rates of photosynthesis and respiration were determined were slightly different from conditions in the culture. There are two effects which could be said to have a bearing on the determination and on the obtained results, namely, total alkalinity and bacterial contamination, c. Total alkalinity The deviation from linear dependence of oxygen supersatu-ration in the culture on rate of net photosynthesis may be attributed to the shaking of the culture prior to rate measure-ment. Shaking thus could introduce an amount of carbon dioxide from the air to the culture changing the condition under which photosynthesis was taking place. This influence may be ob-served in Figure XXVIII, where the level of oxygen supersatura-tion is seen to be inversely proportional to total alkalinity. The highest concentration of Nitzschia closterium in 8 9 these cultures ranged between 10 to 10 c e l l s / l i t e r , in agreement with the maximum obtained by other workers (Ketchum et a l . , 1949; Ketchum and Redfield, 1949). At these concen-trations periodic fluctuations in the number of organisms occur regularly involving birth of new organisms and decomposi-tion of dead ones. Stanburgy (1931) attributes many of these fluctuations and periodic decrease in number to the siliceous shells of the diatoms dissolving in solution at high pH. 100 Aleem (1949) used the method of determining the amount of s i l i c a in the diatom frustules as a measure of their photo-synthetic periodicity. He expressed in terms of unit weight per volume the amount of s i l i c a to indicate the amount of active phytoplankton in a sample of diatoms. It appears, therefore, that close connection exists between the level of pH or total alkalinity and metabolic activity of phytoplankton. This photosynthetic inhibition appears to be more dependent on the effect of the chemical action on the siliceous frustules than on the supply of carbon dioxide. The evidence for this last statement comes from Tailing (1960), who demonstrated this point by supplying bicarbonate to culture water at high pH. He thus showed the maintenance of growth under buffered conditions with the u t i l i z a t i o n of bicarbonate ions for carbon fixation instead of carbon dioxide. It is reasonable to suggest then, that at sea or in laboratory cultures, photo-synthesis may proceed as long as carbonate or bicarbonate ions are supplied in the medium; and inhibition due to pH may occur only when the level reaches a degree where i t could become harmful to the chemical composition of the organisms them-selves. d. Bacterial contamination The presence of bacteria in a unialgal culture is of certain significance in the conditioning of the mode of growth of phytoplankton. In the stationary phase of growth bacteria may accelerate to a large extent the rapid breakdown of the dead organisms and hence supply basic needed requirements to the livi n g c e l l s . This rapid breakdown acts in an opposite manner to photosynthetic activity by lowering the pH through the release of carbon dioxide. On the other hand the oxygen demand for any rate of bacterial growth under both light and dark periods may alter the apparent rates of photosynthesis and respiration. Errors of this nature and their observed influence on measurements and growth have been discussed by Pratt and Berkson (1959). They pointed out for example that during measurement of respiration rate at temperatures ranging between 11-21°C. in two-day light and dark bottle experiment, bacteria were responsible for 42.5% to 62.5% of the total respiration customarily attributed to phytoplankton. In this work, one particular culture, sample no. B^, showed a large concentration of bacteria which appeared under the microscope during c e l l count. The sample results used in Figures XXIV, XXV, and XXVIII are designated differently from the rest to contrast the marked deviation from the other points. Although bacterial contamination is not ruled out completely for the other samples, no further comments in this connection are possible for lack of sufficient observa-tion. 102 A ( i i i ) Relation of catalytic activity to oxygen produc-tion in Chlorella and Nitzschia cultures. The release of cellular compounds into the waters in which phytoplankton grow (Allen, 1914; Gran, 1931a; Matsudaira, 1939; Harvey, 1925 and 1940) has been of prime interest regard-ing their growth stimulative or inhibitive action. Certain properties were found by Pratt (1942) for Chlorellin, the growth inhibitor formed by Chlorella vulgaris. In some cases these substances were found to be thermolabile. Jorgensen (1956) found that Nitzschia could form autotoxic substances while Asterionella species formed substances which accelerated the growth of Asterionella cells and those of Nitzschia, in some cases. Furthermore, he suggested that growth inhibiting substances formed by algae could account for the unialgal nature of a bloom occurring in any natural body of water, where rarely any other species is ever apparent in the v i c i n i t y . For detectable and identifiable compounds in cultures of many organisms, i t has been shown that secretion of extracellular products to the medium parallels that of c e l l multiplication. Guillard and Wangersky (1958) reported parallel increase in number of seven species of phytoplankton to concentration of carbohydrates secreted. It was also demonstrated earlier by Lewin (1956) that production of polysaccharides in C. parvula cultures appeared to be concomitant with growth and proportional to number of organisms. 103 The close association between secretion of extracellular products, phytoplankton growth and variation with catalytic activity was demonstrated successfully by Matsudaira (1950, 1951, 1952a). By noting that the aging effect of cultures, as well as the addition of organic materials similar to those secreted by the organisms, simultaneously inhibited the cata-l y t i c activity of water, he was able to conclude that inhibition of catalytic activity is also accompanied by growth inhibition. Variation of catalytic activity for Chlorella and Nitzschia are given here in Tables XXII and XXIII and plotted in Figures XXIX and XXX. A comparison is made with the level of oxygen concentration reached in the culture six hours after the beginning of light period. The parallelism of both quantities with time of culture for both species is unmistakable. It may prove f r u i t f u l to compare Figure XXIX with that obtained by Guillard and Wangersky (1958) for the production of extra-cellular carbohydrates relative to number of organisms with time - Figure XXXI. Although the diagram is for the flagellate Isochrysis galbana, a similar relationship was found to hold for Nitzschia, where a reported maximum carbohydrate in solution reached 26 mg./l. The close association of supersaturation, increase of catalytic activity and production of extracellular carbohydrate to the increase in number of organisms in the exponential stage of growth reveals that decomposition of hydrogen peroxide is proportional to extracellular product 104 TABLE XXII VARIATION OF OXYGEN CONCENTRATION WITH CATALYTIC ACTIVITY - NITZSCHIA a-x (sample a ( d i s t i l l e d corrected for Kcat. PH water) as pH and S°Q Time °2 of ml. Thio- as ml. Thio- hr.' 1 Days ml./l. sample sulphate sulphate xlO 3 4 7.63 8.10 18.65 17.70 2.19 6 9.50 8.45 18.40 15.30 7.70 11 12.00 8.95 18.13 14.25 10.04 14 15.00 9.20 18.00 13.68 11.44 19 14.50 9.10 17.50 13.95 9.43 23 15.75 9.17 17.32 13.60 10.09 TABLE XXIII VARIATION OF OXYGEN CONCENTRATION WITH CATALYTIC ACTIVITY - CHLORELLA a-x (sample a ( d i s t i l l e d corrected for Kcat. pH water) as pH and S7^ Time °2 of ml. Thio- as ml. Thio- h r . - 1 Days ml./l. sample sulphate sulphate xlO 3 4 7.13 7.95 18.65 17.80 1.96 6 7.50 8.00 18.45 17.05 3.28 11 8.00 8.40 18.15 16.40 4.23 14 9.50 8.70 17.95 15.70 5.57 19 9.50 9.00 17.45 14.75 7.00 23 12.75 9.25 17.25 14.25 7.97 Oxygen measured by micro-gasometric technique secretion and concomitant with growth. The decomposition of hydrogen peroxide in a water medium is dependent, beside other factors, on the chlorinity and pH of the sample, with decomposition increasing with increase of 105 either chlorinity or pH over 7.0. By correcting for these two quantities there remain the action of metal hydroxide ions on the la^^decomposition: Cu (Matsudaira, 1952b), Cr, Fe, Ni, Os, Mn and Co, which are but few of the ions that can p a r t i c i -pate in the catalytic reaction (Schumb et a l . , 1955). Catalytic action on ^ r o m biological compounds may be due to many products of the c e l l ; most important among them is the peroxi-dase enzyme which occurs widely in plants (Eliot, 1932). Other substances such as carotene also have been reported to exhibit per oxidase-like activity (Schumb et a_l., 1955) . In order to present a general picture in this work to represent the various reactions of the catalytic activity and mainly to explain the increase of activity with increase of growth and i t s decrease with culture aging and decay, the following points are suggested: 1. The catalytic activity of a sea water sample depends on chlorinity, pH and metal ions in solution in absence of organic compounds or biological growth. 2. When phytoplankton start growing in the sample, the secretion of extracellular products, or the decomposition of dead c e l l s , produce enzymes which may raise the catalytic activity. The main biological enzymes may be peroxidase and carotene-like substances. 3. The presence of extracellular products may act to enhance growth by providing ready-made complex organic compounds, 106 or may inhibit their own growth and other species' growth with semioxidized compounds or denatured enzymes. 4. With further secretion of organic material further inhibition or stimulation may be possible; however, organic compounds may react with metal ions at a certain pH to form chelate compounds which may precipitate as particulate matter. This step could deprive the organisms of their trace metal ions and inhibit their growth and at the same time prevent the metal ions from acting as catalysts, thus inhibiting the decomposition rate. B. Effect of light and dark periods on oxygen concentration in a heavy culture of Nitzschia closterium The dependence of oxygen concentration on photosynthetic activity, and consequently on duration of light and dark periods as well as on the age of culture, was examined here to account for variation in the saturation level. Table XXIV shows the different concentrations of oxygen at different depths in the culture medium, at three different culture ages, ranging from 22 to 91 days, arid their change with light and dark periods under which the culture was incubated. Each age period demonstrated characteristics sharply different from the others, hence allowing for the discussion of three separate cases: 107 TABLE XXIV VARIATION OF OXYGEN CONCENTRATION UNDER LIGHT AND DARK PERIODS I N NITZSCHIA CULTURE (Oxygen Measured by Winkler and Polarographic Methods) 22 days after inoculation 23 days after inoculation I n i t i a l Dark I n i t i a l Light 2 4 6 2 4 6 8 Depth cm. %°2 hrs. %°2 hrs. 7o°2 hrs. 7o°2 hrs. hrs. 7o°2 7o°2 7o°2 hrs. 7,°2 hrs. 7o°2 2 123 117 113 110 120 121 122 126 128 16 121 115 111 108 118 119 121 124 127 30 119 113 109 Surface 106 Tension 116 118 121 72.8 dynes cm."l 123 126 44 days after inoculation 45 days after inoculation I n i t i a l Dark I n i t i a l Light 5 5 Depth cm. 7o°2 hrs. 7»°2 7P°2 hrs. 7.°2 2 148 145 137 142 12 150 145 137 142 22 151 147 138 140 30 150 Surface 145 Tension 137 72.6 dynes cm.~l 140 90 days after inoculation 91 days after inoculation I n i t i a l Dark I n i t i a l Light 5 5 Depth cm. 7o°2 hrs. 7.°2 7»°2 hrs. 7o°2 2 78 76 71 72 12 83 78 70 75 22 78 76 70 75 30 82 Surface 67 tension 74 71.5 dynes cm.-1 74 108 Case I At the age of 22-23 days from inoculation, the culture was assumed to have just reached the stationary phase of growth. A l l organisms were at the bottom of the container leaving a clear medium for light penetration. No gas bubbles or impuri-ties were noticed either in the body or at the surface of the liquid. Average saturation between dark and light periods was of the order of 1217o during the two-day experiment. Variation in oxygen concentrations at different depths in the tank with dark and light periods followed an exponential decrease and increase (Figures XXXII and XXXIII), with a particular oxygen gradient that accompanied a l l measurements. This gradient, and the fact that oxygen supersaturation was far less than expected in comparison to experiment A ( i ) , indicates a special mechanism for oxygen escape. If in a young culture most of the organisms are expected to be photosynthetically active, and their activity continuous; then the suggestion that microbubbles may result from such activity may be applicable in this case (Ramsey, 1962). The rapid production of oxygen in a closely packed culture produces a strong tendency for the emitted molecular oxygen to coalesce. When this occurs, rapid diffusion into and out of the micro-bubble is possible. The microbubble, being buoyant, starts an ascent to the surface. An oxygen concentration gradient may possibly develop under these conditions i f the aggregate size 109 distribution was a function of the depth of the water column. Otherwise, a concentration gradient may form by the action of diffusion and solubility of the different gases p a r t i c i -pating in microbubble formation. The studies on gas bubbles* st a b i l i t y in liquids by Wyman et a l . (1952) suggest that as bubbles rise and hydrostatic pressure decreases more diffusion of gases could take place from bubble to liquid. Since both oxygen and nitrogen have nearly the same diffusion coefficients, their relative escape from bubble to liquid depends more on their s o l u b i l i t i e s . Oxygen is nearly twice as soluble as nitrogen, thus oxygen diffusion is more rapid to the liquid as the bubble rises than nitrogen. After a certain time, however, the concentration in the liquid becomes limiting for oxygen diffusion and a constant concentration gradient is formed. Case II At the age of 44-45 days from inoculation, the oxygen concentration gradient disappeared, and the average saturation in the culture during the two-day measurements became of the order of 145% saturation. The culture medium became slightly cloudy. The escape mechanism of oxygen is assumed to have reverted to molecular exchange with the atmosphere. Probably, because at this stage oxygen production is more sporadic and intermittent than before, the chance for coalescence due to rapid production is decreased, eliminating the effect of 110 buoyancy and hydrostatic pressure. This allows molecular oxygen to be primarily influenced by molecular and convectional diffusion, and hence remains in the water for a longer time. Case III When the culture reached the age of 90-91 days, the medium became very cloudy, and the brown colour in the organ-isms became more dominant. The surface of the liquid supported a thin layer of impurities composed mainly of bacteria. Surface tension dropped from 72.8 to 71.5 dynes cm.-1, and the average saturation for the two-day measurements f e l l to 75%. It is expected at this time that oxygen production may have been very sporadic and highly dependent, on light penetration through the medium, besides rapid breakdown of organisms and lowering of pH. Bacterial u t i l i z a t i o n of oxygen, and a high demand in the oxidative process of decomposing substances may have been the prime factors in bringing down the oxygen concentration to well below the saturation level. C (i) Influence of water column geometry on rate of desupersaturation To examine for the influence of differently shaped con-tainers, used in previous cultures, on the attained level of supersaturation, the rate of desupersaturation of oxygen was measured for different surface to volume ratios of oxygen supersaturated water. The rate constant Kdes. was a measure of the rate of oxygen escape from physically induced super-I l l saturation in a certain time interval (Findlay and King, 1913). Kdes. - log ^  (17) Where a is the difference between saturation and supersaturation in ml./l. at time zero, and a-x the difference at time t. The results are given in Tables XXV and XXVI for both sea water, salinity 30.20%, and d i s t i l l e d water, and also shown graphically in Figure XXXIV. , The value Kdes. hr."*" i s seen to increase exponentially with increase of surface to volume ratio in both sea and d i s t i l l e d water. Comparison of the rate constant Kdes. for saline water with rate of decrease of oxygen under dark periods in experi-ment B for the same surface:volume ratio, gives further evi-dence to the proposed mechanisms in Cases I and II. The calculated Kdes. (biologically active) for the 22-day-old culture is .107 hr." 1, and .0190 hr." 1 for the 45-day-old culture. While the calculated Kdes. (organisms-free) for the same surface:volume ratio of 0.0323 cm. ^ gives .0186 hr." 1. This is an indication that organism consumption of oxygen in the f i r s t case, i.e., in the younger culture, and i t s proposed escape by microbubble formation exceeds that of molecular escape by about six times. On the other hand, most of the oxygen loss in the older culture (45 days) is due to molecular escape through the surface. 112 TABLE XXV EFFECT OF SURFACE TO VOLUME RATIO ON DESUPERSATURATION RATE CONSTANT t = 12-1°C. S = 30.2%, Surface; Kdes. Volume a a-x t cm. "1 0 2 ml./l. 0 2 ml./l. hrs. 0.0322 6.15 5.30 8.00 0.0417 5.27 4.12 7.00 0.0541 7.30 5.79 7.00 0.0690 7.74 5.80 7.00 0.1053 7.08 4.55 7.50 0.2000 6.94 3.69 7.00 0.2500 7.50 3.40 8.00 hr. -1 0.0186 0.0352 0.0331 0.0412 0.0590 0.0902 0.0986 Surface: TABLE XXVI EFFECT OF SURFACE TO VOLUME RATIO ON DESUPERSATURATION RATE CONSTANT + t = i 2 - r x . D i s t i l l e d Water Volume a a-x t cm.-1 0 2 ml./l. 0 2 ml./l. hrs. 0.0322 3.70 3.28 7.00 0.0400 3.54 3.15 7.25 0.0500 4.30 3.43 7.50 0.0714 4.52 3.57 7.00 0.1111 3.90 2.73 7.00 0.2000 4.10 2.25 7.50 0.2500 3.67 1.82 7.50 Kdes. hr. -1 0.0172 0.0161 0.0302 0.0337 0.0505 0.800 0.935 113 C ( i i ) Influence of salinity on rate of escape of oxygen under stirred conditions The escape of oxygen from stirred or quiescent water is governed by a molecular layer at the surface which i s usually at complete equilibrium with the atmosphere above i t . The escape is governed in quiescent waters by the magnitude of the concentration gradient that forms subject to molecular or convectional diffusion. The effect of stir r i n g normally is to decrease the thickness of the surface layer thus increas-ing the rate of escape, or to increase the concentration gradient resulting in the same effect. Further extreme condi-tioning in the surface layer, such as rapid evaporation or breaking of the surface enhances escape of gas to the atmosphere. The oxygen desupersaturation rate constant was measured under stirred conditions for water ranging in salinity between 0 and 32.82%o. The rate constant Kdes. is shown in Table XXVII and Figure XXXV to increase with salinity. These results are in accordance with investigations of oxygen solubility in d i s t i l l e d and ion-containing water in the f i r s t part of this work. The tendency in ion-containing water is to achieve a more rapid equilibrium with the dissolved gas than in ion-free water due to the weaker bond formation with the ion-containing water. TABLE XXVII VARIATION OF DESUPERSATURATION RATE CONSTANT WITH SALINITY UNDER STIRRED CONDITIONS Kdes. a a-x t hrs. hr. -1 0.00 4.10 8.20 16.41 21.30 24.61 32.82 48.6 45.3 46.7 42.1 44.9 47.3 41.2 18.8 15.4 9.5 14.4 9.4 10.2 11.2 2.0 2.0 2.5 1.5 2.0 2.0 1.5 .475 .561 .637 .715 .782 .767 .869 Oxygen measured by polarographic method C ( i i i ) Influence of diatomaceous earth and surface tension on desupersaturation rate constant The use of diatomaceous earth as a means to examine for nucleation and microbubble formation of gases in the presence of photosynthetic organisms was necessitated by the d i f f i c u l t y of employing li v i n g or dead organisms. This is understandable when photosynthesis, respiration and oxidation are considered in such systems. The errors which may be introduced by these side reactions could completely cover any apparent influence resulting from nucleation. The nearest similar substance that may act in the capacity of diatoms, and provide some information would be the aged deposits of diatoms themselves. Table XXVIII gives values of Kdes. for the surface area of diatomaceous earth added to water under stirred conditions. 115 TABLE XXVIII INFLUENCE OF DIATOMACEOUS EARTH WALL EFFECT ON Kdes. - Stirred -S%„ = 27.70 Temp. = 12-1°C. Diato- Surface maceous Area Kdes. Earth a a-x t gm./l. m 0 2 ml./l. 0 2ml./l. hrs. hr." 1 0.000 0.00 7.45 5.55 7.00 .0420 0.031 1.03 5.75 4.05 7.60 .0468 0.062 2.06 7.40 5.50 7.00 .0423 0.124 4.12 6.47 4.65 7.50 .0440 0.186 .6.18 5.85 4.05 8.00 .0459 0.248 8.24 6.55 4.80 7.00 .0444 0.372 12.36 6.55 4.75 7.00 .0459 TABLE XXIX INFLUENCE OF SURFACE TENSION ON Kdes. - Stirred -S%o= 33.01 Temp. = 12-1°C. Cone, of Surface Heptanoic Tension Kdes. Acid dynes a a-x t mg./l. cm."1 0 2ml./l. 0 2ml./l. hrs. hr." 1 72.9 6.55 4.30 7.00 .0601 50 72.6 7.15 4.65 7.00 .0615 100 72.2 8.00 5.25 7.00 .0601 150 71.6 6.00 3.85 7.00 .0633 200 70.9 4.65 2.80 7.50 .0677 250 68.2 5.97 4.07 7.00 .0640 300 66.8 4.60 3.00 7.00 .0610 350 64.1 4.00 2.70 7.00 .0561 400 61.3 4.30 3.10 7.00 .0546 Oxygen measured by the Winkler method 116 However, as shown in Figure XXXVI, no apparent effect is observable. The results may mean that either these aged organisms lack certain surface substances which could o r i g i -nally have induced bubble formation, or actual bubble forma-tion results only from a highly active photosynthetic region in the presence of the live organisms. Addition of surface active agents, however, proved more f r u i t f u l . The desupersaturation rate constant increased with decreased surface tension upon the addition of small quantities of heptanoic acid then decreased with further addition, Table XXIX. Figure XXXVII shows the change in the rate of desupersaturation with surface tension. These results are similar to those obtained by Eckenfelder and Barnhart (1961) for the influence of surface tension on the coefficient of transfer of a gas, such as oxygen, to the liquid from a gas bubble. They found that the transfer coefficient decreased rapidly with decrease of surface tension, then increased or remained constant on further addition of substances such as peptone and heptanoic acid. The decrease in surface tension on addition of surface active compounds, results from the concentration of these molecules at the surface of the liquid. With small additions, the surface forces decrease and allow for a more rapid escape of the supersaturated gas to the atmosphere. However, further addition produces a monolayer film at the surface of the liquid 117 which may act as a shield for further escape of gas. A large proportion of the gas which escapes from the uppermost thin layer is also dependent on the rate of evaporation of the liquid. The presence of the organic monolayer film may cut evaporation to about 0.01 of i t s value in pure water (Bikerman, 1958, p. 83). Therefore, more addition of surface active agents results in either a constant or a decreased rate of escape. 118 CONCLUSION Experimental results in this work have confirmed the following, previously known points: 1. Oxygen supersaturation in sea water i s directly related to photosynthetic organisms. 2. Oxygen consumption due to respiration is related to the concentrations of carotenoid pigments in the organisms. 3. Oxygen production and photosynthesis in diatoms are influenced by total alkalinity and pH of the medium. On the other hand, the following new relations are deduced from experimental results: 1. The biological activity of Nitzschia closterium in saline water ranging between 25/Uand 31%,, and at temperature of 12°C, exceeds that of Chlorella strain "A". 2. Oxygen production may be related to the difference in concentration between chlorophyll-a and carotenoids under the same light intensity and duration and at constant tempera-ture. 3. Extracellular production of organic compounds and the catalytic activity are concomitant with growth. 4. In new and highly active cultures, microbubble forma-tion may be one of the oxygen paths of escape to the atmosphere. 5. Depth of the water column, stir r i n g and salinity influence the rate of escape of oxygen from the supersaturated water to the a i r . 6. No observable e f f e c t on the rate of oxygen escape can be a t t r i b u t e d to the presence of s i l i c e o u s substances, as demonstrated by addition of diatomaceous earth to the supersaturated water. 7. A small decrease i n surface tension on addition of organic, surface active agents increases the rate of escape of oxygen from the supersaturated water. Further addition of the organic substance decreases surface tension and also decreases the rate of oxygen escape. 120 F i g u r e I . P r o p o s e d D i p o l e O r i e n t a t i o n o f a Water M o l e c u l e i n P r e s e n c e o f I o n s (From Hindman, 1961) 90° Pulse t i mer Pulsed oscillator 1 ® Ho "Sample tube F i g u r e I I ( a ) . NMR A s s e m b l y Insulation Wooden board -3ZZZZZZZZZJ lead Sample Induction coil Copper coi I Thermo couple F i g u r e I I ( b ) . I n d u c t i o n C o i l Chamber To diffusion ft vacuum pump pie tube magnetic bar magnetic st i rre r F i g u r e I I ( c ) . D e g a s s i n g I n s t r u m e n t A ^ — t o gas cylinder Po lye thy lene tube F i g u r e I I ( d ) . Sample Tube 1 A o o A t 1 1 \ V / ~ \ 1 > First Second echo pulse pulse F i g u r e I I I . O s c i l l o s c o p e D i s p l a y o f t h e Two P u l s e s and t h e D e t e c t e d r - f S i g n a l l o g 4 A v s . t i m e to to F i g u r e I V. D e t e r m i n a t i o n o t T, T E M P . ° C . F i g u r e V I . Smoothed C u r v e s o f T, V a l u e s a t D i f f e r e n t T e m p e r a t u r e s ( S a l t S o l u t i o n ) Figure VII. Comparison of T, Values for Oxygen-free D i s t i l l e d Water 127 o ° c 0.| I I I I I I I 0 10 20 30 40 50 Cone , m g. / I • F i g u r e I X . V a r i a t i o n o f 1/T^ V a l u e s w i t h Oxygen C o n c e n t r a t i o n s ( S a l t S o l u t i o n ) T E M P °C F i g u r e - X . B e h a v i o u r o f l / T 1 ( c ) w i t h T e m p e r a t u r e ( D i s t i l l e d W a t er) .25 r \ 1/T. . , vs. temp. C x \ C ) .24 .23 o o> in *i .22 .0.5 M sodium ch lo r ide • 21 r .20 \ N v O 0 10 - o 20 T E M P °C . 30 40 Figure XI. Behaviour of 1/T , N with Temperature (Salt Solution) 1(c) r-1 Figure XII. Change of C ^ Term with Temperature II.0 B co 2 rr  V S o t e m p '  c s ' S ' e 0.5 M. sod ium c h l o r i d e s o l ' n . CM 9.0 7.0 5.0 CD CU 00 -</> 3 D i s t. w a t e r 3.0 1.0 10 20 30 Figure XIII. Change of 40 T E M P . ° C B 50 60 Term with Temperature 2.0 r log ^ c(o) vs. —, T CT> O ro O ro CM 1.0 D i s t. water 0.5 0.5 M sod ium chloride sol'n 0.0 2.8 2.9 3.0 3.1 3.2 1000 . -| . — - — deg. 1 k 3.3 3.4 Figure XIV. Determination of the Activation Energy E( 3.5 3.6 r-1 CO ro .700 log Te(o) v s . ^ 1.300 0.900 Dist. water 0 .500 2.8 2.9 3.0 3.1 1000 3.2 deg . " 1 k 3.3 3.4 F i g u r e XV. Determination of the A c t i v a t i o n Energy E g at D i f f e r e n t Temperatures ( D i s t i l l e d Water) 1.0 1°8 *t*e(o) v s . ~ 0.5 M sod ium c hlor i de so l ' 3.0 3.1 3.2 3.3 1000 3.4 deg . ' k 3.5 3.6 F i g u r e X V I . D e t e r m i n a t i o n o f t h e A c t i v a t i o n E n e r g y E a t D i f f e r e n t T e m p e r a t u r e s ( S a l t S o l u t i o n ) F i g u r e X V I I . D e t e r m i n a t i o n o f t h e H e a t o f S o l u t i o n f o r Oxygen i n Water u n d e r 1 Atm. P r e s s u r e F i g u r e X V I I I . V a r i a t i o n o f t h e R a t i o o f t h e Two R e l a x a t i o n Terms - E q u a t i o n (8) - w i t h T e m p e r a t u r e 137 F i g u r e XIX. Two D i f f e r e n t O r i e n t a t i o n s o f Two D i p o l e M o l e c u l e s 13 CM cO CM O O C J o-/ / o o •_o c ' / ^  "°- O o. \ / \ \ \ N 2 o o -o o P H 10 PH 0 2 SAT. LEVEL-0 10 20 T I M E 25 D A Y S 30 35 40 45 F i g u r e XX. V a r i a t i o n of Oxygen and pH i n N i t z s c h i a C u l t u r e w i t h Time u> oo 13 I I o ' - O O O O O O - o • o -o o -' N. 00 CM o , / \ / / / ' \ \ / P H 10 z o o / p H 0 2 SAT L E V E L 0 15 2 0 T I M E 2 5 D A Y S 3 0 3 5 4 0 4 5 Figure XXI. Variation of Oxygen and pH in Chlorella Culture with Time t o ^0 F i g u r e X X I I . C o m p a r i s o n o f Number o f O r g a n i s m s i n a N i t z s c h i a C u l t u r e w i t h P h o t o s y n t h e t i c P i g m e n t C o n c e n t r a t i o n o F i g u r e X X I I I . C o m p a r i s o n o f Number o f O r g a n i s m s i n a N i t z s c h i a C u l t u r e w i t h R a t e s o f P h o t o s y n t h e s i s a n d R e s p i r a t i o n t-1 •p-I I I I I I I I — 0.0 0.1 0.2 0.3 0.4 0.5 ' 0.6 0.7 0.8 N E T P H O T O S Y N T H E S I S GROSS F i g u r e XXV. B e h a v i o u r o f N e t : G r o s s P h o t o s y n t h e t i c R a t e s i n N i t z s c h i a w i t h R a t i o o f Pigment C o n t e n t F i g u r e X X V I I I . I n f l u e n c e o f T o t a l A l k a l i n i t y on t h e S a t u r a t i o n L e v e l i n a N i t z s c h i a C u l t u r e •P-147 148 8 -ro O o o 0 0 K cat. 5 DAYS 10 AFTER 20 I N0CULATI0N 25 F i g u r e XXX. V a r i a t i o n o f Oxygen and C a t a l y t i c A c t i v i t y i n a C h l o r e l l a C u l t u r e w i t h Time 149 F i g u r e XXXI. E x t r a c e l l u l a r C a r b o h y d r a t e R e l a t i v e t o Number o f O r g a n i s m s f o r t h e F l a g e l l a t e I s o c h r y s i s g a l b a n a ( a f t e r G u i l l a r d , 1958) 130 125 I 2 0 o 2 cm. » 16cm. " - ° 30cm. o ^ o 1 0 0 2 4 6 T I M E - H O U R S F i g u r e X X X I I . Change o f Oxygen C o n c e n t r a t i o n u n d e r I l l u m i n a t i o n i n a N i t z s c h i a C u l t u r e a t D i f f e r e n t D e pths F i g u r e X X X I I I . Change o f Oxygen C o n c e n t r a t i o n u n d e r D a r k i n a N i t z s c h i a C u l t u r e a t D i f f e r e n t D e p t h s F i g u r e XXXIV. V a r i a t i o n o f Oxygen D e s u p e r s a t u r a t i o n R a t e C o n s t a n t w i t h S u r f a c e t o Volume R a t i o o f a Water Column ro 10 __ __ o o 0 10 20 25 30 35 F i g u r e XXXV. V a r i a t i o n o f Oxygen D e s u p e r s a t u r a t i o n R a t e C o n s t a n t w i t h S a l i n i t y u nder S t i r r e d C o n d i t i o n s 5 r CNJ o I I I I I I 0 2 4 6 8 10 12 S U R F A C E A R E A m 2 F i g u r e XXXVI. B e h a v i o u r o f S i l i c e o u s S u r f a c e A r e a w i t h Oxygen D e s u p e r s a t u r a t i o n R a t e C o n s t a n t i-» 4> 7 60 62 64 66 68 SURFACE TENSION dynes cm 70 - I 72 74 Figure XXXVII. Influence of Heptanoic Acid Addition on Oxygen Desupersaturation Rate Constant Ui 15.6 REFERENCES Abragam, A. (1961) "The Principles of Nuclear Magnetism", Clarendon Press, Oxford. Allen, E. J . (1914) J. Mar. Biol. Assoc. U.K., 10, 417. Aleem, A. A. (1949) J. Mar. Biol. Assoc. U.K., 17, 633. Arnon, D. I. (1959) Nature, London, 184, 10. Autry, R. P. and Cole, R. H. (1952) J. Chem. Phys., 20, 1309. Bernal, J . D. and Fowler, R. H. (1933) J. Chem. Phys., 1, 515. Bernheim, R. A., Brown, T. H., Gutowsky, H. S. and Woessner, D. E. (1959) J. Chem. 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Yentsch, C. S. and Scagel, R. F. (1958) Jour. Mar. Res., 17, 567. .25 r \ l/T.., N vs. temp. °C. 1(c) .24 .23 o CO in " 2 2 .21 .20 .0.5 M sodium ch lo r ide X .19 «. o ± — .T*,.Tfc. 10 20 T E M P ° C . 30 40 Figure XI. Behaviour of l/T with Temperature (Salt Solution) l(c ; .25 r \ l / T , , x v s . t e m p . C . 1(c) .24 .23 t> CD M " 2 2 .0.5 M sodium ch lor ide .21 .20 N ' O >. o .19 v o _L — _ o 10 20 T E M P ° C . 30 40 F i g u r e XI, B e h a v i o u r o f l / T w i t h T e m p e r a t u r e ( S a l t S o l u t i o n ) 1(c) .25 r \ vs. temp. C. .24 .23 o " 2 2 V .0.5 M sodium ch lor ide .21 V .20 [• N - O s N N O .19 J_ — o L_ — » ° ** *> 10 20 T E M P ° C . 30 40 Figure XI. Behaviour of 1/T with Temperature (Salt Solution) 1(c) 2 0.0 rr.\ I 5.0 h i • u </> CM * 10.0 5.0 -1 C ^ sec. vs. temp. °C. o s * N«.»^0.5 M sodium chloride sol'n D i s t. Water 10 20 30 40 50 60 70 80 T E M P Figure XII. Change of C ^  Term with Temperature 2.0 r log t c(o) vs. hi' 1.5 o ro O ro (VI 1.0 0.5 • ° ^ v 0.5 M sodium chJoirJide sol'n 0.0 2.8 2.9 3.0 3.1 3.2 1000 . -| . — - — deg. 1 k 3.3 3.4 3.5 3.6 Figure XIV. Determination of the Activation Energy E I .700 h log TeCo) vs. h 1.300 Dist. water 0.900 0 .500 _L 2.8 2.9 3.0 3.1 1000 3.2 d e g " 1 k 3.3 3.4 Figure XV. Determination of the Activation Energy E £ at Different Temperatures ( D i s t i l l e d Water) 3.5 3.6 1.0 0.8 -o o> 0.6 o ro O 0.4 ro 0.2 0.0 3.0 3.1 3.2 log te(o) vs'. ~ 0.5 M sodium chloride sol 'n 3.3 3.4 1000 -| deg. k 3.5 3.6 Figure XVI. Determination of the Activation Energy E at Different Temperatures (Salt Solution) Figure XVII. Determination of the Heat of Solution -^H for Oxygen in Water under 1 Atm. Pressure Figure XVIII. Variation of the Ratio of the Two Relaxation Terms - Equation (8) - with Temperature Figure XIX. Two Different Orientations of Two Dipole Molecule T 1 1 1 1 r Figure XX. Variation of Oxygen and pH in Nitzschia Culture with Time 13 e "I C\J 05 CM O , u z o u • -o 7L -, 0 o o-/ / • • o o o o o o o o No / PH 10 pH ^9 0 2 SAT. LEVEL _L 8 10 15 20 25 TIME - DAYS 30 35 40 45 Figure XXI. Vari a t i o n of Oxygen and pH i n Chlorella Culture with Time Figure XXII. Comparison of Number of Organisms in a Nitzschia Culture with Photosynthetic Pigment Concentration Figure XXIII. Comparison of Number of Organisms in a Nitzschia Culture with Rates of Photosynthesis and Respiration 0 0 0.2 0.4 0.6 0.8 R A T E OF P H O T O S Y N T H E S I S m g C / l / h r . Figure XXIV. Behaviour of Photosynthetic Rate in Nitzschia with the Difference of Pigment Content o 0.0 0.1 Figure XXV, 0.2 0.3 NET 0.4 0.5 PHOTOSYIWTJHESIS 0.6 0.7 0.8 GROSS Behaviour of Net:Gross Photosynthetic Rates in Nitzschia with Ratio of Pigment Content 110 r 90 3 Q. cn E 70 -o c Q) 50 -o a o c 30 -« I I I 0.1 0.2 0.3 R A T E OF RESP IRAT ION mg.C/ l/hr . Figure XXVI. Dependence of Respiration Rate in Nitzschia on Carotenoid Pigment Concentration l 0.4 1.0 Figure XXVII. Relation of Per Cent Saturation to Rates of Photosynthesis and Respiration in Nitzschia 100 NO 120 130 140 150 160 170 % SATURATION Figure XXVIII. Influence of Total Alkalinity on the Saturation Level in a Nitzschia Culture 10 8 I . w -C ro O S 4 0 0g cone. 5 10 15 D A Y S A F T E R I N O C U L A T I O N 20 - 13 15 II CM O 25 Figure XXIX. Variation of Oxygen and Catalytic Activity in a Nitzschia Culture with Time ro O o| 4 2 £ K cat. > 0 , cone. 5 DAYS 10 AFTER 15 I NOCULATION 20 13 CM 9 O 25 Figure XXX. Variation of Oxygen and Catalytic Activity in a Chlorella Culture with Time 10' i '° 6 cn UJ u 10-10 no. of organisms o. / extraceflular Carbo hydrates JL 5 10 15 20 DAYS AFTER INOCULATION 15 10 CO UJ r-< or a > x o CO cr < o d> E 25 Figure XXXI. Extracellular Carbohydrate Relative to Number of Organisms for the Flagellate Isochrysis galbana (after Guillard, 1958) 130 o r-t£ 125 r-< o 2 cm. # 16cm. -°30cm. 120 115 110 TIME — HOURS Figure XXXII. Change of Oxygen Concentration under Illumination in a Nitzschia Culture at Different Depths _i 8 Figure XXXIII. Change of Oxygen Concentration under Dark in a Nitzschia Culture at Different Depths Figure XXXIV. Variation of Oxygen Desupersaturation Rate Constant with Surface to Volume Ratio of a Water Column 10 _____ o — o T 6 V.' x: t n a> T3 2 10 15 20 25 30 35 Figure XXXV. Variation of Oxygen Desupersaturation Rate Constant with Salinity under Stirred Conditions 5 r O a S U R F A C E AREA _j 8 m2 10 12 Figure XXXVI. Behaviour of Siliceous Surface Area with Oxygen Desupersaturation Rate Constant 7 Figure XXXVII. Influence of Heptanoic Acid Addition on Oxygen Desupersaturation Rate Constant lb) lc) Figure I. Proposed Dipole Orientation of a Water Molecule in Presence of Ions (From Hindman, 1961) 90° Pulse timer Pulsed oscillator / 1 ® H 0 ^Sample tube Figure II (a). NMR Assembly Insu latiorv Wooden board -"ZZZZZZZ2ZA lead Sample Induction coil Copper xoi I Thermo couple Figure II (b). Induction Coil Chamber A. C B sample tube { To diffusion S vacuum pump magnetic bar ^-magnetic s t i r r e r Figure II (c). Degassing Instrument A —— to gas cylinder Po lye thy lene tube Figure II (d). Sample Tube A* At J L J L First pulse Second pulse echo Figure III. Oscilloscope Display of the Two Pulses and the Detected r-f Signal l o g ^ A v s . t i m e T I M E IN S ECONDS F i g u r e I V . D e t e r m i n a t i o n o f T, 10.0 r 8.0 r -6.0 h o 4.0 h 2.0 F 0.0 vs. temp. ° C Figure V. Smoothed Curves of Values at Different Temperatures (Distilled Water) . 0 10 20 30 40 T E M P . °C. Figure VI. Smoothed Curves of T, Values at Different Temperatures (Salt Solution) Figure VII. Comparison of T, Values for Oxygen-free D i s t i l l e d Water o°c Cone , m g. C<2 / I -Figure IX. Variation of 1/T^ Values with Oxygen Concentrations (Salt Solution) T E M P ° C Figure X . Behaviour of 1/T, • v with Temperature (Distilled Water) 

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