Free Energy and Other Data for Eight Compounds of Physiological Interest Citation Hlynka, Isydore (1940) Free Energy and Other Data for Eight Compounds of Physiological Interest. Dissertation (Ph.D.), California Institute of Technology. doi:10.7907/x32s-0h08. https://resolver.caltech.edu/CaltechTHESIS:01162025-230014323 Abstract No abstract. Item Type: Thesis (Dissertation (Ph.D.)) Subject Keywords: (Biology) Degree Grantor: California Institute of Technology Division: Biology Major Option: Biology Thesis Availability: Public (worldwide access) Research Advisor(s): Unknown, Unknown Thesis Committee: Unknown, Unknown Defense Date: 1939 Record Number: CaltechTHESIS:01162025-230014323 Persistent URL: https://resolver.caltech.edu/CaltechTHESIS:01162025-230014323 DOI: 10.7907/x32s-0h08 Default Usage Policy: No commercial reproduction, distribution, display or performance rights in this work are provided. ID Code: 16954 Collection: CaltechTHESIS Deposited By: Benjamin Perez Deposited On: 22 Jan 2025 18:44 Last Modified: 22 Jan 2025 18:45 Thesis Files PDF - Final Version See Usage Policy. 45MB Repository Staff Only: item control page

FffEE ENERGY AND OTHER DA TA FOH EIGHT co:i~P OUNDS OF P£1YS.COLOGICAL Il\ITEHEST Thesis by Isydore Hlynka I:::i Partial Fulfillment of the Requirements For the Degree of Doctor of Philosophy California Institute of Technology Pasadena California 1939 I TAB.LE OF CONTENTS · page Introduction l Experimental Methods Solubilities 2 Dissociation Constants 12 Activities 16 Vapor Pressure of Hydra te Systems 17 fa terials 21 Experimen t a l Data Adenine Solubility 24 Hydration 33 Dissociation Constants 35 Hypoxan thine Solubility 39 Dissociation Constants 49 Guanine Solubility 52 Dissociation Constant 61 Xanthine Solubility 64 Hydration 74 Dissociation Constant ?6 II TABLE OF CONTENTS {conJ Uric Acid Page Solubility 81 Dissociation Constant 89 Allantoin Solu·o ili ty 91 Dissociation Const ant 96 Alloxan Solubility 98 Dissociation Con stants 102 Vapor Pressure of Hydrate System 106 Activity Coeffioients 108 Leuoine Solubility 114 Dissooiation Consta nts 120 Summary of Solubility Data 124 Summary of Dissoci a tion Constants l25 Summary of Other Data 126 Derived. Data Heat of Solution 127 Extent of Dissoc iation of Solutes in Saturated Solutions 131 Activities 135 Free Energy of Forma tion of Undissociated Solu·te in St andard Solution 135 III TABLE OP C ONTEW.rs (eon.) Free Ene r gy of Jtormat ion of I ons In Standard Solution 142 Suppl.amen ·l;ary Data 1 47 Equilibrium of Pur ines and Their 149 Discussion 156 Deg:radat,ion Products Summary Refere nc es 1 79 I NTRODUCTION Purines and their related compounds are widely distributed in nature. Among the substances and biological materials whioh oontein purines as their components are chromosomes, viruses, bacteriophages, sperm, cell nuclei, leaf groivth faotors, ooenzymes, tissue extracts and body fluids. In vivo the purines are readiiliy . synthesized even by the simplest organisms. The study of the energetio relationship of these oom:pounds to one another is of interest in indicating whether their synthesis may proceed spontaneously or whether other energy yielding reactions must also take part. The meohanism of synthesis of purines in turn is of interest in its relation to the much wider field of biology indicated by their distribution. The standard free energies of formation of solid crystalline adenine, hypoxanthine, guanine, xanthine, uric acid, allantoin and alloxan have~~~termined by Stiehler and Huffman {1),(2). For the oonsideration of reactions taking place in an aqueous medium it is necessary to know the standard free energies of formation of the various speoies involved. In order to be able to oaloulate the free energies in solution other data are requir ed 1the determination of which oonstitutes the problem of this thesis. -2-- These inolude the determination of solubilities, dissociation constants, activities and vapor pressures. In addition some data on leuoine have been included because the older values are inaccurate due to the contamination of leuoine with 10.ethionine (3). A complete redetermination of the above described physioo-ohemical cons tants has been undertaken. Utilizing the above described data oaloulations have been made of the standard free energy of formation of ea ch of the compounds in its aqueous solution. The equilibrium relations of purines and their degradation products have also been computed. ErflIR I MENTAL METHODS ·rhe experimental methods adopted will be oonsidered first. Since these differ only in detail from one compound to the next a general desoript1on will be given here. The more particular details will be desoribed in oonneotion with the the individual compounds .. 4isoussion ot A brief oonsideration will also be given to the fundamental concepts essential for a clear understanding of the meaning of the various physioo-ohemioal constants determined. SOLUBILI'l'IES Solubility is essentially the measure of an equilibrium constant in a heterogeneous system in which the solid phase is in equilibrium with the dissolved phase. A general reaction may be written as, Solvent t X --,, X -t- Solvent solid ~ dissolved (sat.sol.) the temperature of the reaction being defined. The e~uilibrium constant is therefore, K (Xdissolved) (solvent1sat,~ol),) (Xsolid) (Solvent(liQ.)) But since 1zsolid)1s in its standard state its aotivity is unit y and in a di lute sol ution the activity of the solvent i n t he solution is nearly equal to that ot the pure solvent, we may write K0 Xdiasolved where the subsaript on K0 indioates that this equilibrium oonstant has the dimensions of concentration. The equi l ibrium constant K0 is usually called solubility in dilute solutions. It is to be noticed that the solubility equi- librium does not imply that the solid phase is anhydrous. It speoifioally refers to that phase which is in equilibrium with the dissolved phase. This is of importance where the solid forms a hydrate and will be referred to in that oonnection. The apparatus which was used for the determination of solubilities included solubility tubes, thermostatically controlled water-bath and a rocking devioe tor shaking the solubility tubes, together with weighing bottles, a drying oven eto. The solubility tubes are merely J... shaped tubes such as may be seen in the diagram in Figure 1. When suoh solubility tube is rooked back and forth a thorough agitation of its contents is obtained. These tubes were generally about three quarters full or less to give maximum agitation. Various sizes of solubility tubes ranging from 15 to 250 ml. were used depending on the solubility of the compound and the amount of saturated solution required for the solubility determination. A constant temperature water bath served to maintain the solubility tubes at 25 .t. .03°Cand 50 ±. .05°C( the two temperatures at whioh the solubility dete~minations were made l.• This water bath .was equipped with a moto r driven rocking device which rouke~ baok and forth the solubility tubes. The solubility tubes were clamped to a horizontal bar oonneoted with the rocker. The arrange- ment of the apparatus was such that the lower portion of the solubility tube was immereed in water. The solubilities were determined in the following manner. An excess of a given solid was introduced into a solubility tube which was then about three quarters filled with redistilled water. The tube was then stoppered to prevent oontemin~tion with oarbon dioxide, ammonia etc. This was merely a precaution in order to minimize the sources of er1"'or in the determinations. The solubility tube ,vas then :placed in the thermostat and agitated. Sufficient time was all&wed for the solubility equilibrium to be established. -6- In order to ascertain whether equi librium had been reached solubility determinations were made on several suocessive days. Constant values were ob- tained i11 most cases in one or two days. From supersaturated solutions the equilibrium was in some oases more slowly attainod t han from the undersaturated solutions. Further oonfirmation of the attainment of equilibrium was obtained by ap proaching it from both the undersaturated and the supersaturated side. To a pp roach the equilibrium from the undersaturated side the solid was simply a gitated with water in a solubility tube as already described. To approach the equi librium from the supersaturated side the solid in presence of water was first e quilibrated for one day as described above, the solubili ty tube was removed fra_rn the t h ermost/l).t and hea ted in a water bath for 15 to 20 minutes at 35°0 for the solubility determinations at 25°c, and at 65°c for the solubility at 50°c. The solub1ihi ty tube was then replaced in the thermostat and eq_uili'brated for another one or two days before samp les were removed for determination. For t he solubility determina tions portions of the sa tu.t a ted solution were transf er·red di r ec t ly into a weighing bottle by means of a siphon and suction as shown in the illustration in Figure l. A cotton plug was inserted into the end of the siphon to serve as a filter. In some cases it was neoessary to use good grade Gooch oruoible washed asbestos in order to obtain a elear filtrate. The asbestos was held in place by a small cotton plug on either side. During the transfer of the saturated solution from the solubility tube to the weighing bottle in no oase was there any tendency observed of the so l ute to crystallize out due to sraall tempera ture differences between the solu·tion and the siphon. During the rnanipule.tions the weighing bottles were kept stoppered in order to prevent loss of water by evaporation. The wei ghing bottles used were carefully cleaned in a chromio aoid bath washed and dried in an oven. Under this treatment the lena glass bottles ollanged in weight very little from day to day. Q,uite often their mass remained constant to within .1 mg although over a longer period of time the change in mass was larger than this. All the weighings were done on an analytical balance to .1 mg. A sliglltly larger margin was allowed in determining the 1u.e.ss of the solution. For the gravimetric procedures the mass of the weighing bottles was determined beforehand then samples 'RE.MOV&N<s SOLU.B\ LITY SAMPLES of the saturated solution were taken and the bottles with the solution were weighed. These were placed in a drying oven at 100° 0 and eva.pora ted t o dryness~ During the evaporation the lids were so kept on the weighing bottles as to protect the solutions from contamination with dtt~t particles. At all times the oven was used only for quantitative work . Under tt1ese conditions the evaporation proc eeded smoothly without bumping and consequently with no loss of solution and \Vi th no contamination. The residue was dried to cons ta.nt vveigl'lt also a\ 1000c . Several preliminary experiments drying the residue at 11000 and at 135°c showed that there was no noti c eable ohange in the wei ght of the res :t due at the hi gher temperature. Drying the residue at 100°c was al so used by Dalton and Sohmidt in their solubility determ1.nations of the amino aoids (4). The mass of the residue of the solute was obtained as the difference between the mass of the weighing bottle plus the residue and its mass empty. Similarly the mass of the solvent is the difference between the weighing bottle plus the residue. Due to the faot that most of the compounds were only slightly soluble the mass of the residue was small, oft.he order of magnitude of 10 mg. Since this was determined as the difference between two large numbers a high acouracy was required in the weighings. Thus .1 mg difference in the mass represents a 1% error. Suoh faotors as the absorption of a thin film of moisture on the relatively large surface of the weighing bottle a.re s ignif ioant in this type of work. In order to minimize the errors due to the ohange in the mass of the weighing bottles themselves parallel experiments using distilled water were done with. eaoh determination. The sec ond wei ghing bottle was then used as a counterpoise or tare in wei ghing the weighing bottle plus the residue. rhe mts.sses of the empty 1 wei ghing bottles were also determined by the counterpoise method. For t his ;iurpose the weighing bottles were paired in such a way that their masses were within .1 gm.of eaoh other. In this way the wei ghings oould be done using only the ohain of the oha.inomatio bal ance and were therefore independent of the brass weights. Furthermore sinoe eaah bottle was nearly of the same .dimensions and mass and had undergone the same treatment in drying any changes in their masses would be nearly identical and would thus tend to oanoel out in the counterpoise weighing. -10- For the least soluble members of the purine group naruely uric acid, xanthine and guanine the gravirnetrio method wa.s not suff:toiently accurate so that oolorimetrio methods were adoptedo The a.mounts of :x:ant 11ine and guanine in solu•i;ion were determined using the Koessler and Hanke diazo method (5). The diazotized sulfa~ilio acid in the reagent aouples with xantlline and guanine in an all{a- line medium to yield azo dyes. The oolor that is developed can be measured quantitatively in a colorimeter. The application ·:) f this method to a que.n•!ii tative determination of purines was made by ITunter (6). Since solutions of ~ure compounds were used the usual objections of the non speoiftcity of the reaction were absent. Furthermore suitable standards as will be later described treated in the srune way as the test solutions were used for comparison. Oare was taken to make the two solutions to be compared of nearly the same concentrations. In this way the several errors of the method were minimized. All the solubilities were done on at least two different preparations of each compound where sufficient material was available. The temperatures used were 25 ± .o3°o and 50 ± .05°c . . The solubilities were expressed as grams of solute per kilogram of water. The results of the solubility determinations for any one oompound were conveniently tabulated, the arrangement in the table being as follows. One t able was rnade to inolude the solubilities obtained when the solubility equilibrium was approaohed from the undersaturated side. A seoond table included the results when the equilibrium was approached from the supersaturated side. In the first column of the table the determinations are numbered consecutively. The second column gives the length of time during which the solubility tube was equilibrated before the determination was made. are given in the third oolumn. The solubilities The fourth oolumn represents the deviations of the individual results from the group mean. The group mean was obtained by averaging the results of the solubility determinations . for any one preparation of the compound under consideration and is given at the end of the group in the third column. The final oolumn gives the deviations from the general mean. The general mean was obtained by avera gin g all the results in the table. An analysis of the results in this manner brings out more olearly -12- the variations in the different preparations of the same compound and also their relation to the final result.. The precision errors were calculated a ccording to the method of Rossini (7) using the formula :t 2Yu2. /n.(n-1) 2 where "E..4 is the sum of the squares of the deviations from the mean and n is the number of determinations~ In order to obtain the final values of the solubilities the results approached from t he undersaturated and t he supersatur ated side of the e quilibrium were ave raged . The :precision 2 errors were combined using the formula ! -Ya2 + b , where a and bare the respective precisi on errors of the t wo individual results. DISSOC IATIQ:r~ CONSTANTS In all oases except one the dissoc iation constants we r e determined by eleotrometric titration using a glass electrode and a Beckman pH meter t o follow the H + oono~ntration~ The pH meter was fi:rst adjusted using a phthalate buffer of known pH of 3.97. Then a given solution of known volume and concentration was titrated with a standard a cid or base. 'l'he concentration of the solutions titrated was determined by the solubility of the compound -13- in question as for example a solution of xanthine which was .0001 molar was used on account of the sparing solubility of xanthine in water. The oonaentrati on of the standard acid or alkali used ln tlle ti trations was in turn determined by the conoentra tion of the solution titrated. The use of a micro burette of l ml,oapacity which could be easily read to .005 ml. made it possible to use the standard reagent of higher concentrations thus avoiding dilution of an already dilute solution. In this manner maximum accura c y was obtained from the ti tra ti ons . The ti tratio:ns were done by adding known amounts of the standard acid or base to a given volume of the solution being tit.rated the pH of the solution being recorded after ea.ch addition of the reagent. Care was taken to have the solutions at 25°0 during the titration. The aocuraoy of the pH determinations under the usual working oondi tio.ns was about ,1.05 pH uni ts. In the hi gher alkaline range the glass electrode becomes inaocurate. A correction was applied from a graph supplied with the instrument for this purpose. All the titrations were done in duplicate and the results were averaged. The volume of the reagent added was plotted against the pH of the solution to give a -14- titratio~ curve. Si~ee the dissoeiatien constant varies with the dilutie~ a correoti0n is desirable . A titrati0n of the water blank was therefore done with each titration using the same volume of the solution (water) and of the reagent. The valume of the reagent used for the water blank was subtracted from the a.mount used in titrating the e.quec:nis acid or base fer each :pH thus giving a corrected titration curve . The corrected. tii:ra.ti0n curve thus representl;;I a graphic method of ebta1ning a disaociatien curve . Suoh a eorreeted titration curve is s0metimes used te ebta.in tlae dissociation constant directly from the graph (8) . It is pref~rable,however, te calculate the dissociation constants from the titration data by mes.E.S of appropriate equations. The dissooiatien constants were calculated frem t~e equations and -15- where cA is the concentration in moles per liter of all forms of the substanoe titrated. CNa+ 1s the total conoentratidm, of Na + in moles per liter 001- is the total concentration of Cl- in moles per liter. Cgt iS the oanoentration of g t in moles per liter pKw is 14.00 at 25°c (9). Guanine was found to be too insoluble to be titrated in the above manner henoe the solubility method of Hitc hoook (10) was used. Xanthine was done by both the titr~metrio and the solubility methods. The relation between solubility and dissociation is given by and where, sis the solubility of all forms of the substance in grams per liter at a given pH. S0 is the solubility of the undissociated substanoe -16- whioh is usually taken as the solubility in water. It is also expr essed in gr ams per liter CH+, ~ • Kb, Ka have the same significance as before. The results obtained from the titrations were averaged to give a final value. Sometimes the first and last values were not included in the mean, as the aocuraoy ot the dissociation constants is least when the ratio of the free aoid to the salt is far from unity. The solubility method however can only be ap plied when the ratio of the f!"ee aoid or ba se to its salt is quite small. It has been found to be fairly accurate under these oonditlon~o ACTIVITIES At infinite dilution dissolved substances behave as perfect soluteso At hi gh er concentrations they deviate from the laws of perfect solutions and a correction factor is required to make them conform to these laws. This taotor is known as the aotivity coefficient and the oonoentrations aorreoted in this manner are known as activities. The relation between oonoentration and activity may be expressed by the eq_uation M.'( = a where M is the oonoentration in molality, '( is the activity ooefficient and a is the activity. --17-- In pra.otioe the activity coefficients are determined only for the more soluble substances. As already stated above in dilute .solutions the laws of perfect solutions are more nearly obeyed hence it is common practice to assume as a gretty good a pproximation that the activities in dilute solutions are equal to concentrations. The line of demarcation is usually made at .l molal oonoentra.tion where the activity correction is already small and where the experimental methods begin to be inaoourate. The molal solu.bill ties of the purines we r e found to be less than .01 and that for allantoin .0373. It is therefore safe to assume their activities eq_ual to their concentrations. The molal solubility of leuoine was found to be .16 '/ just on the borderline where the activity measur~ments become inaccurate. Alloxan alone was found to be sufficiently soluble for an experimental determination of its activity ooeffioients. VA POR PHE::3SU'~}£S OF' HYDRATE SYST!lll!S Just as in the case of solubilities the hydration of substanoes oan best be considered from the point of view of heterogenous equilibria. Let us consider a substance X: mich forms a hydrate X. H..,9 0 . Tl1is hydra te will de- compose to form water and the anhydrous s ubstanee X. If such a system is put in a closed vessel an equilibrium -18- will finally be esta blished X .H O ( S) ~ H O ~ '""' 2 2 cg.pJ ' .....anhydrous(a) where pis the equilibrium p r essure of the hydrateanhydrous transformation. The equilibrium oonstant for this reaction is given by K . (g,p~ x Xanhydrous(s) H20 Now sinoe X anhydrous and X, H20 are in their standard states their activities are unity. The activity or fugaoity of wa ter is equal to its vapor pressure, hence the equilibeium oonstant becomes Kn .i; =p where p is a gain the e quilibrium p r essure of the hydrateanhydrous transformation. 'rhe subsori ..nt of Kp indicates that t he equilibrium oonstant has the dimensions of pressure., Let u s suppose a c a se wl1ere ::{ forms a trihydrate but no intermedi a te h ydrates. The e quilibrium involved is then given by X.3H20 U:1\-;;:::! 3H2o (g, p ) -t X anhydrous{s) Tile equili briu.111 constant is therefore K ~ 3 x Xanhydrous (s) X.3H20 {s) and Kp -- n3 s: It will be seen that i t is the equili brium constant tha t is of interest for the purpose of free energy aon- siderations and not the vapor pressure although in the oa s e of a monohydrate these are identical. It shoul d also be po i nte d out that the vapor pressure of t he hydra te does not refer to t he hydrate alone but to tile e q_ui l i bri um syst em• .A.mong the compounds studied alloxan , adenine and xanthine are known to fo rm hydra teso I t was neoessary to determine whether these hydr ates were fo:i.~med at 25°c an d if stable at this temperatur e it was also necessary to de t ermine t he v apor pre~sure of the hydrate-anhydrous $ys tem . The other comp ounds studied did not form hydrates. The genera l method employed will now be consi dered. Wher e thi s was fea s ibl e the sol id phase was completely dis so lve d in water at eleva t ed te mpera tur es. It was then allowed to come to equili bri um by agitating it for several days in a solubility tube pla ced at 25°0 in a thermostat. A sol ubility determination was made in order to asoertain whether the equilibrium had been attained. The solid phase was then r emoved by filtration on a Hirsoh funnel and superfici a lly dried between fi&ter papers. This solid -20- was them :placed on a ta.red. watch glas$ and its weight determinei . It was then placed in a small vacuum chamber together with about 5 ml. of sulfuric acid solution oontained in a small glass vessel. The system was evacuated on the water pump for one or two minutes and finally set away in a 2.5°c incubator . After several days the sample was removed and the loss in weight determined . The rationale of the method is the following . The sulfuric acid eelution acts as a desiccant t0 remove a.n.y water of wetting that still adhe ~es t0 the erysta.ls after superficially drying the s0lid . If the vapor pressure ef the hydrate system i.~ less th.an that ef the au.l:furic ac id solution no further desiccation will take place after the water of wetting has been removed . If the solid is then dri ed in an aven at 100°c the water present will be re- moved . The water removed in this w,anner must then have been bound as water of crystallizat ion. On the other hand if the vapor pressure of the hydrate system is greater tho that of the sulfuric acid solution all of the water will be lest to the desiccant . By using a series of sulfuric acid solutions 0£ varyi~g concentrations the vapor pressure of the hydrate system calll. be determined from the vapor pressure of that sulfuric a.eid solution to which the hydrate neither loses ner gains any water . Experi mentally this is obtained as the average of two ooncentratlons nesr one ano"ther, to one of which the hydx·ate system loses water vapor and not to the other. The vapor pressures of the sulfu~io a.old soluti.olis of' different eoneentrations were obtained from the International Cri tioal 1l'ables (~.l. ) ~ The oonoentra.tions were cheoked by specifio gra'1ity determinations-. These . were done by weighing 5ml,;of the solution on the balanoe. Ml\. 1r ERl ALS A brief account of the s ,:m roe and the method of prepa ration and 1.:,urifioation of the corn.p<:nt(l.ds used in this investigation is g iven below. A number of these were t he preparations described. by Stiehler and Huttman (l} and they will b e desi gt1ated as combustion s ai.11ples. Adenine I - oornbustion s ample , adenine a. Adenine II - Tl11s was prepared by the hydrolysis of yeast nucleic aeid by the method ot Hunter and Hlyuka. (~ ) . The adenin~ hydrochloride obtained was orystal- lized three tim.es from dilute hydrochloric · aoid. It was then conver·ted i n to the free base by dissolving the .hydrochloride in water and neutralizing the solution with sodium carbonate. The free adenine whioh p:cecipi- tated out was crystallized tour times from water . Miorosaopio bipyramid crysta ls were obtained • .Attempts to grow larger crystals ;were unsuccessful as adenine deposited as a orusty preoipitate on the walls of the container. •rhe needle crystal fo:rm as in Adenine I above oould not be obtained under any of the conditions tried which included seeding adeni ne s,)lutions with these crystals and orystallizing at different tempera ... tures and varying dilutions . Guanine I and C1anine II were both combustion samples. i:iypoxanthi11eU, II, and III were al.so combustion smaple s .. Xanthine I we.s a c,.,mbusti 0n s ample. Xanthine II and III wBre Hoffm.e.nn La Roo he prepara- tions decolorized with chiarcoal an.d crystallized four times from water. Xant.1.Line lV was a Hoffmann La Roche preparation treated the se.me way as II and III above. t-, specia l effort v1as made t o obtain l a r ge crystRls by first hea. ting ~a.nth i nc i n _presence of a large amo unt of' its s a turated solution to oo 0 c and then cooling it at e0 c over night in a refri gerat or. This :proc edure was repeate.d for a period of one week. Urio Aoid I was a combustion sarr~le. Uria Acid II was Pfanstiehl C.P. urio aoid orystallized t hree times from water. It crystallized readily to yield well defined erystals • .Allan toin I was a Hoffmann La Roob.e preparation. Allanto1n Il was a comtmation sample. Alloxan was a Hoffmann La Roohe preparation Qryste.1lized onee from water and dried in vacuo over sulfurio aoid. d- and. 1- Leuoine were methionine free samples prepared by Fox ( J. ~) • -24- ADENI NE THE SOLUB I LITY OF ADENINE The solubility of adenine was determined by the gravimetric method . .App roximatel y 10 gram portions of the saturated solution were used for each determina- tion and the wei gh t of the residue obta ined on evaporation was determined as already described. The results of the solubility determinations of adenine at 25°c are given in t a bles I and II. Taking the mean of the general means we obtain for the sol ub ility of c'lden ine at this temperature 1. 13 t .04 grams :per 1000 grams of wa ter. The va riati on of the ind ividual determinat ions constitutes the largest error. The variation of the results on the two diffe r ent preparations is not . sfgni:ficant. Also there is lii. ttle dif fe rence between the values obtained app r .9 ?ohing the e q_uilibrium fro m the undersatura ted or the supersaturated side. It also apper1rs tba t satura tion is obt a ined in a relatively short ti me from the unders atura t ed side but t ba t t here is a ·tendency for the so lution to remain supers a turated \Vh en the eq uilibrium is apuroaahed f rom the s 1.roe r s a turated side.· ~ ~ -26- The results of solubility deter.min.a ti ons at 50°c are . given in tables III and IV. The results obtained give for the solubility of adenine at this temperature 3.4 ! .2 grams per 1000 grams of water . It is at onoe apparent that there is some disagreement between t h e i ndividual resul ts on the two different prepa rations of adenine. Adenine I and Aden ine II as will be reoal1ed differ in orystalline form and it is likely that their solubilities differ, this difference beaoming more apparent a t higher temperatures. Unfortunately an insufficient amount of Ade nine I was a vailable. As has been mentioned above no suooess was ha d in a ttemp ting to prepar e the needle cryst ::1 lli:qe form of adenine. The discrepancy between the i ndividual determinations is not large and the agreement between the values obtained when the equi l ibrium was approached from the undersaturated and the supersa turated side is also quite satisfac t ory. According to Koessel ( .J.~ • 92 grams of e.denine dissolve i n 1000 gr ams of wat er at room tempe rature. Tafel and Aoh ( .15) found its sol.lbili ty in boiling wa ter to be 25 grams p e r 1000 grams of wa te r· . These da t a are of limited quantitat ive value only, but i nJi c u te t h e s ar:ie order o f ma gnitude. "· - 26- It is convenient·te have the solubilities expressed in terms ef m0lalities. At 25°c the molality ef a saturated solution ef adenine from our results was f0und t0 be . 00836 and at 500c . Oj52 For the purpose ef calculating t he percentage dissociation the pH of a saturated solutien is required. It was found to be 6.6 at 25°c . -27- TABLE I SOLUBI LITY OF ADENI NE at 25°c Equilibrium Approached from the Undersa. turated Side . Ex.• No. Days Equi- Sol'y Dev.from Mean Dev.from librated Gm/Kg. Water l l 1.. 17 .03 .• 04 2 w l l .~21 .07 .08 3 w ½ 1.. 12 -.02 -.01 4 ~ 1.10 -.o4 - .03 5 l 1 . 16 .02 .03 6 2 1 . 06 -.08 -.07 7 3 1,. .• 13 -.01 8 4 1.13 -.01 .oo .oo g 1 1.15 .01 .02 10 2 1 . 16 .02 .03 11 3 1.22 .08 .09 12 4 1.21 .07 . ·08 13 2 1.,07 -.07 -.06 14 3 1 . .09 -.05 - ·. 04 15 4 1.14 .oo .01 16 3 1.11 -.03 -.02 17 4 1.15 .01 .02 18 5 -1.21 1.14 .07 .08 Gr. Gen Mean ADENI NE I :t-.02 -28- TABLE I Exa No. (con) Days Eq_uilibrated Sol 9 y Gm/Kg. Dev.from Gr. Mean Dev .from Gen.Mean Water ·. ADENINE I I 19 2 lell .04 -.02 20 3 1 . 10 .03 -.03 21 4 1 . 05 -.02 -.08 22 2 1 .04 -.03 -.09 23 4 1.05 -.02 -.08 24 5 - 1.09 1 . 0? .02 . -.04 :r -:15'2" Mean of group means 1 . 10 :c . Oj General mean 1.13 ;:t • 02 W after the experiment number i n the first column indicates that water was added to the solute of the pre vious determination . - 2 9- TABLE II -0 SOLUBILITY of ADElfINE at 26 C Eq_uilibrium Approaohed from the Supersaturated Side Ex. No. Days Equilibrated Sol'y Gm/ Kg. Water Dev. from Dev.from Gr. Mean Gen. Mean l 2 l.33 -. 19 .20 2 w 2 1 -. 18 .04 3 5 1 -. l? -.03 4 w 5 1 ,. 14 .• oo 5 6 1 . 10 -.04 6 7 1 ,. 10 .:...o4 .;. .• 03 7 w 7 1.08 -.06 -.05 8 6 1 .• 04 - .. 10 -'o09 9 9 - .06 -.05 10 10 1.08 . 1 -. 11 11 w 6 1.17 ·~ .Ot ADEIH:NE I ADENINE II -.OJ l..,14 :t •0 .5 12 3 1.14 .01 .01 13 5 1.18 .05 .05 14 6 1.12 -.01 -.01 15 6 1.14 .01 .01 16 8 1.14 . 01 .01 - 30 - TAaLE ExJNo. II ( con) Days Equilibrated Sol'Y Gm/ Kg. Devoi'rom Gr. Mea.n Dev. !'rom Gen. Mean 17 8 1.11 - ·. 02 -.02 18 18 1.11 .... 02 -.02 19 10 h!Q. -.03 1.13 -.03 :t .02 Mean of group means 1.13 :t .o-., Genera l mean l.13 :t .03 Water -31- TABLE I II SOLUBILITY of ADENI NE at 50°0 Equilibr ium Approached from the Undersa ture. ted Side Ex. No .• Days Jlqui- librated Sol'y Gm/Kg Dev.from Gr. Mean Dev.from Gen. Mean Water ADEN i f\TE I l l 3.18 .oo - .16 2 2 3.22 .04 .:. .12 3 4 3.14 -.04 -.20 4 5 3 .1? -.01 't. .o~ -. 17 3:T8 5 1 3.45 .03 . ll 6 l 3 . 47 .05 .13 7 2 3 . 41 -.01 .07 8 2 3.44 .02 .10 9 3 3.36 -.06 .02 10 3 3.42 .oo .08 11 4 3.38 .... 04 .04 12 4 ') Ol;j .10 ADENI NE II 3.44 3 .42 - :t .03 Mean of group means 3.30 ;- . 04 General mean 3.34 .J: • Ot3 • - 32- Ti\.BLE IV SOLUB !LITY of ADE1'U NE a t 50°0 H:q_111libr1 um Appro aoh ed f r om t h e Aup ersaturated 8 ide Da ys Equi- l ibr e.ted i)ev . i'rom Gm/ Kg . 1 fo ter Dev . from Gr . Menn Gen . Mean Sol' y AD"SN I N.E I l 2 ~ . 46 . 2,2 . 04 ,:~ •"l 3 :5 . 34 . 10 -. 08 ;j 4 -23 .. 2491 - . ~;3 -. 51 ::t- 33 A. DIN I NE II 4 l 3 . 68 • 19 . %6 5 l 3 . 77 •c'.b . 35 6 ( .;J 3 . 39 -.10 -. 03 •) ? ,,., •') 3 . 5g - • I0 -. ,, 8 3 3 . 41, -. Ob -. Ol 9 3 ~ . 44 -. o., . 02 10 4 ::s . 4;) -. 06 . 01 11 4 3 . 40 - . 02 o. 49""- :" . o~, ± • L)'d :; . 36 :t ._54 Genera l mean • 3 . 42 .:t . 1:, Mean of group means (.) .) -33- Ad en i ne is lmown to form both a monohyd:ra te and a trihydra te. I n ord er to determ i n e whethe r t he solid p ha s e .i n e quilibri w.n. with a s a tu1·a ted so.Lution of adenine at 25°C is hydra ted or not the following experiments were d one. Ade n ine wa s completely dissolved in hot water and e q uilibrated in the s a me wa y R s in the solubility experiments. After four d ays e 111 j_libriu..r11 wa s attained as judged by a solub ility determina tion. was then removed 1:':i nd The solid phase su;_J erfici a lly dried b e t ween filter pa pers. It should be possibl e to remove the water of wetting by drying the s a mp l es of aden ine, prepared a s above, in p r es ence of a c.~.sioaant the vap or pressure of wh ioh is hi gher t han that of t he hydr a te system but lower tha n tha t of t he sf:l.t u:cs.t ed solu tion t)f ad enine. }txper i ments of t h is t ype were J.one usi ng s ul furic a cid s olutions of known vapor pressur es. In t a ble V below1 the va por pressures in mm. of Hg . of the sulfuric aoid solution used is g iven i n t h e seco nd co.lumn. In th e third, fou r th and fift h columns a r e given t he wei ghts of the s up erfici a lly dried sample, of the s ame s amp l e dri ed f or four du ys in presence of a s :1 lfu.ric a ci d so l u t ion a n d t he wei ght of the sample dried in a n oven at 100°0. The last column -34- indic a tes the numb er of molecules of wa ter rema ining a s w~ter of hydration . All experiments were done at 25 °c. This series of e xperi me nts indicates ttla t aden i ne does not form a hydra te under these conditions. The last exper i ment in t he seri es s hows tha t wa ter of weting is not remove d when the superf'ici a l ly dried s amp le of adenin e is :p l a o t~d in p resence of wa ter alone. Al- though some wa ter appears t o have been lost i t was shown to be due to a prelimina ry loss during the evacuation of the va cuum chamber. Even if a hydrate were formed having a vapor pressure of t his order of magnitude it would contri bute little to the free energy chang e and co uld be ne g l ected. TABLE V HYDRATION OF ADENI NE AT 25°c No Vap. Press of 'V t of Super- 1 Wt of Desi cc. Wt of Moleo of Dried 4den. a q . H2so 4 Soln. fio. Dried Adenine Dried Adenine l 19 .40 . 229 . 186 .18 6 0 2 21.40 .09 1 .067 .067 0 3 22 .54 . 095 .07 5 .075 0 4 23.25 .09,8 .073 .073 0 5 Wa ter . 148 .128 • .103 1.8 oven .... 35_ DISS OCIATION CONSTANT S OF .ADENINE AT 25°c The dissociation constants of adenine were determined by eleatrometria titration. Fifty ml. of .00222 molar adenine solution was titrated with 1.2 ml. of .0989 N HCL and Na OH respectively. Each titration was done in du~lioate and the pH readings we r e averaged. The dis- sociation constants were c a lculated as described above. A titra tion of a water blank was also done. are given in tables VI a n d VII. The results Figure II shows the titration c urves of t lle adenine so..i..utions a nd the correspondin g water clanks . A corrected titration curve is also given. An average of the pK values included between asteriaks was t a ken. The value obtsined for uKa is 9.90 whic h gives for ~ the aoid dissociation c : x:i.s tan t Ka a value of 1.26:xlO -10 The value of pK-0 was found to be 9.86 which g ives for the basic dissoc i a t ion consta nt Kb a value of l. 3 8 x: .J.O -10 • Tl1e li mit in6 i'ao ·tor in t he ace uraoy of the above titra tions is t h e high dilutio ns of the solutions which 1 t wa s necessary to employ. The constanqY. of the dis- sociation c onstant in the middle range of the titration -36- curve is a n i ndicati~n that the ~alues obtai ned are quite satisfactory. ,.2. ADENINE TITRATION CURVE.S APE.NINE. o- o- o WAT&I\ BLAr.U~ "- - ' " x.- "'- C.OR R!C.,-£ .::P CURV~ - - - - - 8 10 u- · J: ·3 -. J: 0 :t: ~ z ~ l:.3 4- pH 6 - 37- TABLE VI ACID DISSOCIATI ON CONSTANT OF ADENINE AT 2 5°c NaOH pH of Ade11ine So.l.ution pH of Wa t er Blank pKa o.o 6 . 57 6. 50 •••• .1 8 . 76 10 . 08 ',J . 77 .2 9 . 25 10 . 68 9 . 89* .3 9.48 11 . 00 9 . 92 .4 9 . 67 ••••• 9 . 92 .5 9 . 83 11 . 3 ? 9 . 93 .6 9.';39 ·• ... . 9 . 93 .7 10 . 14 ... .. .8 10 . 28 .9 10 . 41 .. . . 1.0 10 . 55 11 . 85 ~.. 62 l .l 10 . 71 •• • •• s , ga 1.2 10 . 87 ml ~ , ~3 . .. ... 9.88 9 .80 * .. . . . •••• Average pKa Ka l . 26x10 - 10 -38- VII TAB.LE BASIC DISS OCIATI ON CO WS'rANT OF ADENI NE AT 25 ° 0 HCL pH of Adeni ne Solution pH of Water Blank pKb o.o 6 . 31 5 . 83 . .. . .l 5.22 3 . 68 IJ.81 .2 4 . 85 3.35 9.83* .3 4.62 j . 18 9 . 84 .4 4.44 j . 04 9 . 84 .5 4.29 2 . 96 9 . 85 .6 4 . 14 2.86 9 . 86 .? 4 . 00 2 .82 9 . 86 .8 3 . 86 2.74 'd . 85 .9 3 . 72 2 . 69 9 . 88 1.0 3 . 58 2.64 9.89* 1.1 3.43 •••• 9.94 1.2 3 . ~o .... .. . . ml. Average pK b Kb 9 . 86 , -10 l.' ·tlxlO - 39- HYPOXANTHINE THE SOLU]3ILITY OF BYP OXANTIIINJ:~ Hyp oxa nthine solubility determ. i na tior1s we re done by the gravimetric method. I)o rtions of' about 20ml. of the s a turated solu tion were requi r ed to g ive a r esidue of a bout 14 mg. Since the c apa cit y of the we i ghing bottles used was about 10 ml. t wo successive s amples wer e evap orated to dryness an d the total resid ,..1e wa s de te rmi ned. Tl1e results a re given in t ab les VIII to XI. At 25°0 t he solubility was f ound to be .73 1: .01 grams per 1000 gr a ms of wa te:r. 1, l tho ugh all the prepara- tions were ca re fully purified combus t ion s amples there app ei:1.rs to be so me v a ria tion in their so.Lubi.Li ties. It is .nore l.i.!rnly that t hi s v ar iation is due to the low prec i s i on of t he me thod. In a ddition it also appears t ha t ap_p roa chi ns t he e q uilibrium fr om t;tae supersaturated side the solutio ns tend to rema in supersaturated for several days. At 50°c the so lubi lity was fo und to be 2. 01 -+ grams per 1000 g r a ms of wa ter. li t . 06 hi gher temp e ra tures hypoxanthine is more soluble and the gravi metric method gives resu lts of g re a ter precision. TJ1e agreement among the individual determina tions as well as among the different preparations is muc h better. 'l'he tlg.ce emen t between the results obta ined from the unders a tura ted solution ~nd those from the supers a t urated solution is a lso bet t er at the h i r4l er temper-a.tur e. T.he so lubi lity of hypoxan thine in water has been dete rm ined by several inv es tigB. to.cs. A summary of their results is given for compar ison. Temper·ature Solub ility in Gm/ Kg Water Investigator (16) Ro om 1 . 03 Bruhns " .63 " Tt . 53 " .g 2 " Scherer l 17,} 11° c 1.50 Stutzer (18 ) 19°c .71 Fischer ( J.9) 23°c . 73 l't 100°0 14.30 " •rhe results obtained by Fischer are in s ubstant i al a greement with ours. For convenience our re s ults s. r e given here in terms of molali ties. The solu.bi l i ty o f hypoxan thine at 25°c was found to be . 00536 moles per 1000 grams of wa ter and 0 at 5o c .0148 moles per li ter. The pl:l of a s a tur a ted solu tion wa s fo und to be 6.4 at 25°c. - 41- TABLE VIII SOLUBI.LITY OF' HMl-,OXA N'rliINE AT 25°c Equilibrium Approached from t he Undersaturated S ide Ex ,. No. .Days 3 q ui- So l 'y libra ted Gm/ Kg . Dev. fro m Gr . Mean Dev.from Gen. Mean ~1va t er HYPO- XAFTHI NE I 1 2 .'684 -.031 -. 0.54 ~ 3 .681 -.034 - • 03'j 3 4 .716 .00 1 - .OC£:: 4 5 .787 .0?2 • 06<)· 5 w 2 .666 -. 049 -.052 6 4 .'740 .025 . 022 ? 5 .'710 -. 005 -.008 8 5 .758 -• 04,3 .040 g 7 .7?6 .o•6;'1 .050 10 w 3 .702 - .0 13 -.o b 11 4 .694 -.021 -.0~4 12 5 .710 -.005 -.ood 13 6 .769 .054 .0 5'1 14 ? .7 04 -.011 -. 0 14 15 4 .695 - . 020 -.Oc:.3 16 5 .693 -.022 -. 025 l? 6 • 697. -.fU8 - . 02 ·1 18 7 . 69.l -.0~4 -.027 8 -.• 715 -06 -.009 - • r 1 C. l 9 7 -t . m r, -42- TABLE Ex. No. VIII (con) Dev. from Gr. Mean Dev.from Gen. Mean .682 -.025 -.036 3 .705 -.002 - • 0 :1 3 B2 4 • 729 . Oi: :'. 2 . 011 23 5 .703 -0 004 -.0 :15 24 6 .701 -.006 - • 0 l '.J 25 w 3 . 701 -.006 - . 01 :, 26 5 .721 .014 . OOj 27 6 .'743 .036 • ()'.c) 28 w 3 -.681 .707 -.026 -t: • Ofl' t - • o:;/J 29 2 .735 -.00 1 • 01 i 30 3 .748 .012 • Q2J) 31 4 .72'7 -.009 .009 32 5 .740 . 004 .022 33 6 .'751 .0 15 .033 2 • 722 -.0 14 • 004 35 3 .733 -. 003 .015 36 4 . . 7Jl ~ -.005 .013 .:t Days Equi- ~3ol' y librated Gm/Kg . Water 20 2 21 HYPS - XANTHI NE II HYPO- XA NTHI NE III 34 w Mean of group means General mean .7 19 .718 .oo'f :t.0,2t:: ::t . 010 TABLJE IX 3 0 .LillH LITY OF HYp OXAJVi Hn rn t ~ra.J. 25°c 1 Eq ut .Librium Approache d from the Sup ersa tura t ed L~ ide Ex. N'oQ Jays TE'.,tU1 - ~3ol 'y Gm/Kg Dev . from librated Gr . Mean Dev . from Gen . Mean t'/a ter HYPOXAXTHI !m I l 2 . 750 . 018 .ooo ...,,_;, 3 . 739 ., Q07 - . 011 ~ 4 . '725 -. 007 - .025 4 5 . 741 . 009 -. 009 5 6 . ?05 -. 02'7 -.045 6 7 . 718 - . 014 -.032 7 8 . 720 - .• 0.12 -.03 0 8 9 • '736 .0 04 - ,.014 9 10 . 711 - .021 - .039 10 11 - .045 . OB 7 . 777 . 7:32 :t . oa -4- HYPO- XA NTHINE II 11 3 .748 - .o~ -. 00 2 12 4 ·• 740 - .0 10 -. 010 13 5 .7 59 . 009 .oog 14 8 .745 00 5 - .005 15 ~ .?49 -. 001 - .00 1 16 10 ~ .750 . 010 .0 10 <AO ':l' . 006' -44- TABLE IX ( con) Days J!.:qui- Sol'y libra ted Gm/Kg. Dev .from Gr. Mean Dev .from Gen. Mean Water XANTHHJE III 17 4 18 5 19 6 20 0 '768 • 0 .1.8 .005 .024 .7'12 .003 .022 7 . ,.?83 , '·· .014 .033 21 8 ·808 .039 .058 22 w 3 -751 -. 018 .001 23 4 .739 -.0;30 -.011 24 5 -.006 .013 25 6 761 ... od3 .011 26 7 •768 -0001 .± . o1; .018 . 769 Mean of group means .750 :± •019 General Mean .750 ± .009 0 ,r - 45 - TABLE X 0 SOitJBil.ITY OF HYPOXANTHlNE AT 5o o Equilibrium Approached from the Undersa tu:ea ted Si de Ex. No. Days Equilibrated Sol'y Gm/Kg . Dev .fr om Dev .from Gr . Mean Gen. Mean Water HYPO- XANTHI . E I 1 1 2 .02 .0 2 .0 1 2 2 1 . 'J 6 -. 04 -. 0.5 3 3 i . 00 .oo - • o ·i 4 4 l. 9'/ -003 -.04 5 B 2 . 0 .5 ~.oo .0 5 .± .03 .04 6 6 1. 9 4 - .02 - . o·, 7 1 2 . 05 . 09 .04 8 2 1. 9'7 .O l -. 0 4 9 3 2 .00 .0 4 -.01 10 4 1.82 -.14 - • l~ 11 5 1.98 .o i:'. -.03 12 6 1 . 9'/ • O· i :t .0 5 -.C4 HYPOXA:'iTHI N:E II l.96·. -46- TABLE "Ex. No . X (oon) Days Eg_ui- Sol'y librated Gm/Kg. Water Dev.from Gr. Mean Dev.from Gen.Mean HYPOXAN"TITI NE III 13 1 2 .17 .11 .16 14 2 2 . 18 012 • 1.7 151 3 2 . 01 -.0) .oo 16 4 2.0 c -. 04 .0 1 17 5 2.01 -.05 .oo 18 6 1.98· 2.06 - • Q 't j -.0.3 + .0:1 Mean of group means 2.01 -t- . 09 General mean 2.01 :t . 0 4 -4?- TABLE XI SOLUBILITY OF HYPOXANTHIN~ AT 50°c Equi librium Approached from the Supersaturated S ide Ex. No. Days Equilibrat ed HYPO- Dev .from Gm/Kg. Water Gro Dev.from Mean Gen. Mean Sol'y XANTHINE I 1 2 1.91 -.Ob -.13 2 3 2.01 ,04 -.03 3 4 1 ~96 -.O J -.08 4 5 :i.. . 99 .02 -.05 5 6 1.96 -.01 -.Ob 6 7 o ffi(; . r ....:..9..,6 -.01 j: .03 HYPO- XANTIUNE II 7 2 1.98 -.04 -.06 8 3 2 .14 . le .j 0 9 4 2.04 .oc aOO 10 5 1.99 - • 03 -.0 5 ll 8 1.99. -.03 ·- . 0 5 12 7 2.01 2.02 -.0 1 -.03 ± .05 XI { oon) 'fABLE Deys Equi - • l.i bra te<i S ol ' y Gu1/I(g. G.r.. Mean Dev . from Dev . 1'1:om 2 ., l3 ,.00 .09 Wa ter Gen . Mean HYPOXA:?ffHINE I I I 13 l4 3 2 .. 15 .oc . 11 15 4 2.1~3 , 00 . 09 16 5 2 . 10 -.03 . 06 17 e 2 . 14 • C·1 • I0 18 ? 2 , 15 2.15 . o~ ± .o~ Mean of group means 2 .04 .:i General mean 2. 0 4 ± . 04 . 06 .1, -49- THE DISSOCIAT ION CONSTANTS OF B.TPOXANTHI NE AT 25°c Both the acid and basic dissociation constants were determined by means of an eleotrometrio titration. Fifty ml. of .0022 molar hypoxanthine was titrated with 1.2 ml. of .0989 N. HCL. The same quantity of this solution was also titrated with l.2 ml. of .09r-39 M. MaOH. The results obtained are given in tables XII and XIII. These are rep.resented graphically in Figure 3. The average value for pKa was found to be 8. 97 whioh gives for the acid dissociation constant of hypoxanthine a value of l.0?xl0 -9 • The variation of pKa values is s mel l except for the last few. The basic dissociation constant of hy-poxanthine is quit e s mall. Due to the faot that it was n ecessary to use very dilute soJ..utions of hypoxanthine for the titrations a nd also beoause the dissociation constant is small aacura te results could not be obtained. However taking an average of the pKb values in the middle portion of the titration ourve some idea is obt a i ned as to the order of magnitude of the basic dissociation. The PI% value ob- tained is 12.? and hence t he dissociation oonstant oomes out to be in the neighborhood of 2.0xlO -13 0 :B'ili tti (20) using the solubility method found the acid -, dissociation 12 constant of hypo:x:anthine to be 2.12xlo- .Iti. a Sl.lbsequent :paper she reported Ka 2.13x10•-"ll and tb 1.L0'Jx10- 1 .3 at 400c. -50- TABLE XII THE ACID DISSOCIATION CONSTA:'JT OF HYPOXANTlU NE A'f 25°0 pH of Hypox. Sol'u of pH Water Blank •o 5.50 6.50 •••• .l 7.99 10.08 9.01 .2 0.35 10.68 9 .01 .3 8.59 11.00 9.02 .4 8.7? ••••• 9.02 .5 8.93 11.31 9 .02 .e 9.08 .. . . . 9 .01 .7 9.23 ••••• 9 .00 .8 9.38 ..... 8.9? .9 9.54 ••••• 8. 91 1.0 9.15 11.85 8.81 1.1 9.99 •• • • • 8.91 1.2 10.30 •• • • • •••• NaOH ml Average pKa Ka 8.97 l.07xl0 -9 ,.z. TITRATION CURVES 0 f HYPOXANTHINE. ~ ~~ijTtllfr\le. ·9 \1A-i"C f\ el.AN'< ~Oft~e.CTE. D o-o - 0 -D ~-~-- "- c.u~ve. _ __ _ • . 0:i::·3 . \. 0 ·9 2 4 pH 10 -51- 'J.1ABLE XI II THE BAS IC DI SS OCL\ r I ON CONST.ANT OF HY?OXAN'rHI N.t.!i Arr 25°0 HCL pH of Hypox. Sol•~ pH of Water Blank ,:pKb 6.12 5.82 ••••• 3. 78 - 3.68 12.06 3.45 3.35 12.29 .3 3.26 3.18 12.56* .4 3 ol2 3 .04 12 . 73 .5 3.02 2.96 12.93 .6 2.95 2 .86 12.66 .7 2.89 2.82 12 . 4 5 .8 2.82 2.74 12.87 .9 2.78 2.69 12 . 58,.: 1.0 2.72 2 .6 4 13.07 1. 1 2.68 •••• 1.2 2 .62 •••• Average pK0 Kb 13.04 0 12 .7 2. C- xlO - 13 *Average t a.ken of values between asterisks. • • • • -52- GUANINE SOLUBILITY OF GUANI NE For the quantitative determina tion of guanine in s a tura ted solution a colorimetric method was employed using t he Koessler and Hanke diazo reagent (5). proeedure was as follows. The A standard solu tion of guanine wa s prepared by dissolving 3 mg . of guanine in a liter of redistilled water. The r a te of solution is slow and can be i ncreased by pulverizing the soli d with a blunt stirring rod and he ati ng on a water bath . • Having t he standard, a suitable portion of about 5 ml. of a saturated guanine solution was then withdrawn from a solubility tube. The determi na tion of the amount of g uanine in the s a tur ated solution was then made by ma king the follo wi ng solutions and compar ing them in a colorimeter . a) 5ml. standard solution d.iazo reagent. b) + .lgin. Ne.2003 + 5ml. satura ted guanine sol·ution diazo reag ent. 5ml. +- .1gm. . Na 2 co 3 -+ 5ml. A brown ish oolor develops first and pers ists for a.bout 15 minutes. The solutions f '.lnally become olear yellow and. ar.·e :ready for comp a rison i n about 25 minutes. The initial brownish color is probably due to the slow deamination of guanine. In maki ng a colorimetrio -53- oomparison oare should be t aken to have the same shade of color in both the standard and test solution. It each solution is treated i n the s ame way and sufficient time is allowed for the oolo r to develop this difficulty can in large measure be avoided. An aver-d.ge of 10 read- i n gs on the colorimeter was taken. For the solubility at 50°c the proportion of the rea gents wa s altered as follows. The standard was made by dissolving 5 mg. of guanine in a liter of 1.1% Na 2 co 3 solution. The solutions compared in the color- imeter were, a) ·b) 5ml. sta ndard guanine sol.ution + 5ml. diazo reagent. 2ml. saturated gua nine s ol 1.1tion Na2co 3 solution r 3 ml. of 1 .1% t 5ml. diazo reagent. Ot her details are the same as before. The results obtained for the solub ility of guan ine at 25°c are g iven in tables XIV and XV . The difference between t h e res ult s on two different ,reparations as well as betwee n those obtained from the unders a tura ted and the sup e r saturated si_de of the e q_uilibrium app ears to be with i n t he pr ecisioa of t he method. Combining t he results a s p reviously exp la ined we obta in as the solu- o bility of guan ine a t 25 C . 0034 .± .0001 grams per 1000 grams of wa ter. This corresponds to a. . 000022 .5 molt:>,l solution at saturation. -54- The .0 131 8 ol ·1.1b i l i ty of guanine at 50°0 wac found to be ± . 000 :4 gr ams per 1000 grams of wa t e r. rr he mol a.lit y of t he s a tura ted solution comes out to be .00008 6 '7 at this te m_p ere. tu r e. He r e, as above, the li miting f a ctor in the pr ecision of t he result is the precision of the method. Wood ( 21} deterra..i ned t .he solubility of guani n e at 0 40.l C usi ng t he gravimetrio me th ~Jd and fo und it to be .039 gr ams per 1000 grams of wa ter. Severa l obvious rea sons fo r t h e l ar ge discrepanc y be tw een t he res ults of Wood and our res:..i.lts ma y be of f ered . Sinoe gua nine is not readily crysta l l ized it has to be purified by successive rep reoi pit at io ns . The pur it y of t h e t wo prepara ti ons mi gh t ha ve been diff e r e nt. With compounds having a small solubility wi de ly di ffer ent results mi gh t be obta i n ed if a s ma l l amount of i mpu rity we r e pr esent. A second a.nd a lso a likely exp la na tion mi ght be t he d i ff e r en ce i n the me thods. 'l'he souraes of error i n a gravi metric method in t h e determina tion of the solubili t ies of sli gh t ly s ol1.1ble comp oun ds are oonsiderab l e. It was fo r this re ason t ha t a oo lorirnetric method was a a.op ted in our de t ermina tions .· Further reasons for t he diso r e "iJancy between the r es ul t s obta i ned by Wood and our r es~lts will be suggested in connection with the determ inationation of t he dissociati on c onsta nts. -55- TABLE XIV SOLUBILITY OF GUANH·ill AT 25°c E~uilibrium Approached from the Undersaturated Side Ex. No. Days Ey_uilibrated Sol'y Mg/ Kg Water Dev .from G:r. !Jean Dev .. from Gen. Mean 1 2 3.8 .4 .4 2 2 3.4 3 3 3,4 .o .o ~o .o 4 3 3 • vJ7. .... l -.1 5 4 3.3 -.1 -.1 6 4 3 .4 .o .o ? 5 3. 5 .1 .1 8 5 GUANI NE I ~ -.1 .1 j7I "' ·~ 3 .? .2 .3 10 2 3.5 .o .l 11 3 3.3 -.2 -.1 12 3 3.,4 -.1 .o 13 4 3.5 .o .1 14 4 3.6 .1 .2 15 5 3.2 -.3 -.2 16 5 3.4 -. 1 .o 3. 3 GUANI NE II 9 :'375 -t--:T Mean of grm.ip means 3.4 j: . 1 General mean 3.4 ::t .1 -56- TABLE rv SOLUBILITY OF GUAN I NE AT 25°c Equilibrium. Approached from the Supersaturated Side Ex . No. Days Equi- librat ed Sol' y Mg/Kg Dev.f rom Dev.from Gr. Mean Gen . Mean Wat er GUANI NE I l 2 3.2 -.3 -.2 2 2 3.3 -.2 -.1 3 3 3.3 -.2 -.1 4 3 3.3 -.2 -. 1 5 4 3. 6 .1 .2 6 4 3.6 .1 ~2 7 8 3.6 .1 .2 8 8 ~ ..,Cl, .5 . .3 .4 ::! .2 GUANI 1TE II 9 2 5.3 .... l -.1 10 2 3. 5 .1 .1 11 3 3.2 -.2 -.2 12 3 3.3 -.1 -. 1 13 4 3.3 -.1 -. 1 14 4 ~? . 3 -. 1 -.1 15 8 3.6 .2 .2 16 8 -33 .4. 7 .3 .3 ± . ·; Mean of group means 3.4 ::t • ~ General mean 3.4 -t . l -57- TABLE XVI SOLUBil..I'rY OF GUANINE """" 50°c .~ rri Equilibrium Approached from the Undersaturet ed S ide Mg/ Kg Dev ofrom librs.ted Gr . Mean Dev .from Gen. Mean 1 7 11.• 8 - 1 .2 -1.5 9 ,_., 7 12 .• 0 -leO -1. 3 3 a 12 . 4 -. 6 -. 9 4 8 12.8 -.2 -. 5 5 9 15.0 .o -. 3 6 9 12.5 - o5 -.8 7 10 1;?, , 3 -.7 - 1 .0 13 l 14.4 1. 4 1.1 9 1 13 .? .? .4 10 l 13.6 .e .3 11 1 13.6 .6 .3 12 l 13.2 .2 -. 1 13 2 13.3 .v r,i: 0 14 2 13.1 .1 - .2 16 2 13.0 .o - .5 16 2 13 . 1 • .l . -.2 17 2 13 .$ 13.0 .i .3 E:x. ~No. Jays Equi- Sol'y Wa ter GUANINE I - .3 0 -58- XVI {con) TABLE Da ys Equilibrated Sol'y Mg/Kg Wa ter pev.from Gr. Mean Dev . from Gen. Mean 18 1 12.3 -1.3 -1.0 19 l 13.6 .o .3 20 1 13.7 .1 .4 21 l 13.4 -.2 .l 22 1 13 .2 - .4 -.1 23 2 14 . 5 .9 1.2 24 2 14 . 0 . .4 .7 25 2 13 .4 - .. 2 .1 Ex. No. GUANI NE II 26 i:: ,, 13 .0 -.6 -.3 27 2 13 .4 -.2 .1 28 3 12.5 -1.l -.8 29 3 14 .0 .4 .7 30 3 14.l .5 .8 31 3 14.2 .6 32 3 - 14 . 2 13 . 6 .e i .3 of group means 1~3 .3 j: .4 General mean 13 . 3 + ~? Mean. 0 • .9 .9 -59- TABLE XVII SOLUBILITY OF GUA~i!INE A'.r 50 0 C Equilibrium. Approached from the Supersaturated Side Days Equilibrated Sol'y Mg/Kg Water Dev.from Gr. Mean Dev.from Gen. Mean 1 1 13 .5 .6 .7 2 2 13.0 .1 .2 3 13.1 4 l 13.l .2 5 1 13 ·. 4 .5 0 1 13.0 .1 7 l 13.0 .1 8 1 13.2 .3 9 2 13.5 .6 10 2 12.0 11 2 12.9 12 2 12.5 13 2 12.5 -.4 -.3 14 2 12.4 -.5 -.4 15 3 13.0 .1 .2 16 3 12.9 .o .1 17 3 12.8 -.1 0 18 3 12.8 -.1 0 19 3 12.8 -. 1 0 20 3 - -.l 0 Ex. No. GUANINI? I 12.8 .3 .2 .7 -.8 .o .l -.3 -60- 'fAB.LJ!: E:x . Uo. :..-VII Days Equi- libra ted (con) Sol•y Mg/Kg Water Dev.from Gr. Mean Dev.from Gen. Mean GUA NHffi II 21 j. .L4 .o 1 .4 1. 2 22 l .l,.;:;. 0 , .- -.1 - •V 23 l 12 .7 .l -.1 24 l 12 . 3 -.3 - .5 25 l 12.7 .1 -. 1 26 2 13 .7 1.1 '•g 27 l:3 12.4 -.2 .4 28 t:., 12.4 -.2 .4 29 2 12 .6 .o -.2 30 2 13.0 .4 .2 31 4 12 .8 ' .2 0 32 4 12.0 - •Q -.8 33 4 J.1.5 -1.1 -1.3 34 4 11.6 -1.0 -1.2 35 4 - 13.0 1 2.6 .1 .2 :t j Mean of group means 12.8 j .4 General mean lB . 8 j: .2 'J ,., •) r~ THE BA'.:.HO DISSOCIATION OOMSrr;t:N'r AT 25°c .As has been already po inted out the dissooia tion constan t of guanine was done by the solubility method . . Sme.11 amoun ts of g u.8.nine were put into a series of solubility tubes and equ ilibrated with successively increasing oonoentrations of hydroohlorie aoid for a period of t wo days. P ortions of the saturated solution were removed and the solubility determined as previously d.escribed. The standard solution of guanine used for compar ison was made up i n hydrochloric aoid solution of the same concentration as the solution being tested in order to promote soluti\)n and to have the same conditions for the development of color with the diazo rea gent. S imultaneously with the solubility determina- tions of pH were made. The dissociation consta nts were then c a lo~lated in the manner aescribed above. The value obtained for the solubility of g uanine in water was used as t he minimum solubi+ity. The results obta ined for t h e dissociation co nstant of ~·uan i ne at 25°0 are g iven i n t able XV III. The error in the pH determinations is estima ted as .05 pH units. Since a difference between the minimum solubili ·ty and the solubility in the hydrochlor ic acid is involved in t h e calculation a large error is introduced where this differenoe is s ma ll, ie., in dilute hydroohlorio aoid solutions. Furthermore, the dissociation constant is small and is therefore in a reg ion where its determination is difficult-. rrhe results obtained therefore, show wide variation. The average value of pKb taken between the figures marked wi th asterisks is 10.90. sooiation constant Kb l •.3xl0 This gives as the dis-11 . This value of the dissociation const an t is taken to represent the order of magnltude rather than an aoeurate result. The variation of solubility of guanine in hydrochloric acid solutions and consequently its dissociation is represented graphically in Figure 4. Wood (2:h) using his solubility values in tenth normal and twentieth normal hyq1·ochloric acid found the 0 basic dis sociation constant of guan ine at 40.2 C to be .807xl0 -11 • The dissociation const ant appea rs to be smaller even at a h:i gher t emperature than that obtained in our experiments. ation. It is su~gestive of some contamin- VARIATION °f • f SOLUBILITY 0 GUANINE with pH 20 lO O - - - - - - - J . - - - - - - 1 . . . . . _ __ 5 4 . ___., _ _____._ _ _~ - - - - - - : pH 3 -63- TABLE XVIII 0 DISSOCIA'rION CONSTANT OF GUANI NE AT 25 0 Ex, No. Solubility pH of Guanine Sol'n in HCL pKb l 3.96 4.57 10.22 2 4.03 4.5? 10.08 3 4.42 4.19 10 . 27 4 4.65 3.'18 10.25 5 4.37 3.69 10.88* 6 4.20 3.39 10.93 7 4.85 3.39 10.99 8 5.00 3,39 10.94 9 6.05 3.08 11.03 10 5.75 3,08 11.08 11 8.?3 2.79 11.01 12 8.?0 2.?6 11.04 13 16.70 2.58 10 .82 14 38.00 2.31 10.68 15 71.00 2.06 10 . 71 16 129.00 1.79 10. 54:ai: Mg/Kg of Water P¾ 10.90 Kb -11 1.3 :telO -64- XANTHINE THE SOLUBILITY OF XANTHINE The solubility of xanthine was determined by the eolorimetrio method of :Koeasler and Hanke (5}. method was as follo ws. The A standard solution was pre- pared by dissolving 3 mg . of finely pulverized xanthine in 200 ml. of water thus giving a oonoentrat1on of 15 ro~. per kilogram of water. Under ordinary conditions the rate of solution is slow but with the aid of the blunt end of a stirring rod and sometil:ies heat solution is finally effected. A s 11itable :p ortion of the saturnted solution of xanthi ne was tro.nsferred from a solubility tube into a wei gh ing bottle. The t ests were performed by adding to each of. the test tubes 2 ml . of t he Koessler and Hanke diazo reagent, followed by 5 ml. of 1.1% Na 2co3 solution to one and an equal quantity of a satura ted xanthine solution to the other. About 20 minutes was allowed for the yellow oolor to develop. Comparison was made in a co lorimeter in the usual way. 0 The quantities of the reagents for the 50 C solubility were m~dified so as to obtain a suitable intensity of color in each of' t he solu tions. The sta ndard wa s made by dissolving 3 mg . of xanthine i n 60 ml. of 1.1% Na 2oo~ - 65- solution t hus giving a concentration of 60 mg . per 1000 grams of wa ter. a) The test was done by using , 2ml. of diazo reagent + 5 ml . of 1.1% Na 2oo~ ~ solution + l ml. of saturated xanthine solution. b) 2 ml . of diazo reagent -t- 4 ml. of lsl% Na 2co 3 solution -t 1 ml. standard solution + 1 ml. water. The solu tions were compared after 20 minutes. The solubility of xanthi ne a t 25°c was found to be .0 164 .:t .0J0 '. / grams per 1000 grams of wat er . A satura- ted solution is therefore t a ken as .000108 mola l. From tables XIX and XX it may be seen that the agreement of the results on the t wo preparations of xanthine used is satisfac tory. There als o app ea rs to be little di ffer- ence whet her the so~ubility equ i l bri1-1m is approached from the unders a tura ted or the supe rsa turated side. The largest va r iation is among the individual determinations. It therefor e a pp ears tba t the precision of the result is det ermin ed b y the p recision of t he r,1ethod. The results of the solubilit y determina tions at 50°c are s hown i n t ables XXI and XXII. The so lub ility of xanthine at t h is temp erature was found to be .oeq .:t .004 grams per 1000 grams of water. The saturated solu tion 1s therefore .000440 molal at this tem:.9eratur e. Although there is a oons iderable var ia ·tion among the ind ividual determinations the l a r gest discrepancy is due t o the fact that xant hine I has a lower solubility than the other t wo preparations. In t his connecti on it is i nstruo ·ti ve to e:xamine t he graph in Fi gure f• If it is reca lled that the solubility of xanthine i n wa ter was fou nd to be 16 . 4 mg. per 1000 grams and tha t it increa ses ver y rap idly with , ,,. a sli ~;tlt ohange in :pH, it will be seen that the so l u'bili ty of xa n 't hine will be considerably aff'eoted by small amou1 ts of cont aminants . In this li ght it is possible to understand the l arge vari a tion in the r esults obtain ed by various i nvesti ga tors. 1~ slilmmary of their results is given below. Temperature Solubiltty Gm/Kg Water Investigator 10°c . 470 Strecker (22 16 .615 16 . 0675 Sundwick ,2 ~S: ) 16 .oeg Almen (~4 .) 17 2.600 40 .075 S taedeler ( 25 ) 40 . 182 \: ood {21 , } 100 1.380 100 .725 100 . 8c!5 S taedeler \25) 100 .??O Almen ( .24) n S tutz.er p.a., Strec ker ~ 2 ) (· n The oonolu~1on t ru.~t haa oeen d1'tra.n from r@sul ts su.-oh i~ a above is tht& t t h e s.o lubili ty of' mnth ine de- pends on the met l1od of prepa :ration. I t is mox-e logical to oonc.lud~ that t .tle S('> iubili ty depe:ndS> rJn the pu:ei ·ey of t h e p:repa1•atio:ns, tUlrl tne eil>ove result$ twiy th.en 'be exp l a ined b y t hf; va :ryi n g ~U1l.{)u11t:s of :lJ;.1pu.ri thls i.n the :ic~nth ine p rej!?t.it'a t ions. rr11e ei:"feot of _ ; purU'ioa t1on is to dEH'J r@e i,;a t h e solti b :L .li ty ot XtH'itl1i:ne. 1·e h®l.1erve 0 tha t o·,.ll" results !'I.re raore ~HHH1ra ta than a n~· of t hose reported i n the li te;r;-u t ure . Our re$ ul ta on the tiis - soo t a t 1on 01.Jnst a nt ot x11nthine rd.so Sll1,lfmrt tt1.is uonC.iU$ion . TABLE XIX SOLUBI LITY OF XA 1J'!'HI N:!!: AT 25°C Eq uili b:rium ap:-proaohed from the Unclersaturated Side Dev.from Dev.from. Gr. Mean Gen. Mean 16.0 -. .:, -.3 2 15.9 -. 3 -.4 .J 11 1 5 .l .. 1.1 -1.2 4 12 17.8 1.6 1. 5 5 14 15.8 -.4 -.o~- 6 16 15.7 -. 5 -.6 7 18 17.4 16.2 1.2 1.1 1: .7 Ex. No. Days Equi ... S ol'y libra ted Mg/ Kg Water l l 2 X ti.NTHI NE I ,.., ,··. XANTHI NE IV 8 2 17.5 1.1 .1 . 2 g 3 15.9 -. 5 -.4 10 5 ,~ r,· l o.o -.9 -.8 11 6 16.5 .1 .2 12 ? .!§..& .4 .oF 16.4 1- • 7 Mean of group m.eans 16.3 +,,.o General mean 16.3 --t- .B TAB.i.,E XX 0 S OLUB I L ITY OF 7..Al:VI'H I NE AT 25 C Ey_uili .bri um App roached from the Sup ersatura ted S ide Ex. No. Equilibrated De. ys Sol'y- Mg/Kg Water XA!fi'If I NE I Dev.from Dev.from Gr. Mean Gen. Mean l 3 17 .1 .3 .6 2 4 18.5 l.? 2.0 3 5 16.5 -. 3 .o 4 6 , 16.3 -.5 -.2 5 15 1'7.5 .7 1.0 6 16 16.1 -.7 -.4 7 17 ~ 16.8 -. :r -.6 G .~ 15. 9 -.1 -. 6 3 15. 9 -.1 -.6 10 5 15.9 -.l -.6 11 6 16.5 .5 .o 12 7 - XANTHI NE -.9 7 IV 8 g - 15.7 16.0 ..... 3 .::t .3 Mean of' gro up rueans 16.4 :t .8 General me an 16.5 :t .5 -.8 -70- TAB.LE XXI 0 SO.LiJIHLITY OF Xt:-.UTHI NE AT 60 G Equi li'br1um Approac h ed. frani t he Unders a tura ted Side Da.ys Equi- S ol'y Devofrom Dev .from librated Mg/ Kg Water Or . Mean Gen. Me~n l l ?0.5 l.4 3.9 2 2 73 . 5 4 .4 6.9 3 3 78.0 8.9 11.4 4 ~:1 70. 2 1.1 ;3 . 6 5 5 00 . 2 -8.9 -6.4 6 7 66. 2 -2. 9 -.4 ? 8 66 .6 - B.5 .o 8 9 58 . 2 -0.9 1.6 Q 10 68. 2 69,l -0.9 ::t 5.0 .L.o 10 l 68 .7 8.9 2.1 ll 2 62 .l 2.3 -4.5 12 3 58.4 -1. 4 b.2 13 4 55 .6 - 4 .2 -11.0 14 6 59.0 -.8 -·,. 6 15 '7 58.4 -1.4 -8. 2 16 8 ~ ~- 6 6' - -10.0 Ex . No. XANTHHfE II XANTHI NE· I 59 .. 8 - 3 .2 _j-J.O -?l- 'l'Ji.B.LE XXI Ex oNo. Days Equilibrated {oon) Hol'y Mg/Kg Gr.. it ean Dev.from Gen. Mean Dev,fro111 Water X..t1. N11 HINE I I I l? l 70 . 3 1 •1 3 .7 .l.8 l '71.4 2.2 4.8 19 2 68.7 -.5 2 .1 20 2 71.4 2.2 4.8 21 3 69.2 .o 2 .6 22 3 6'7 . 7 -1.5 1.1 23 4 69.2 .o 2 .6 24 4 66 . ? -2.5 .1 25 5 68.7 -.5 2.1 26 5 .filh2. -.5 :t .9 2.1 Mean of group means 68.0 General mean 6 b .b ± 5.0 :t. 2 .o 69.2 -72- TABl1E XXII 0 SOLUBILITY OF Xi\ i'.'lJ.~HINE AT 50 C Equilibr ium Approached from. the Supersatura ted S ide Dev.from Gr~ Mean Dev.from Gen . Mean 6'9.? .6 2 .9 3 7 4-.4 2.3 4.6 3 4 68.7 -.4 l. 9 4 6 68.2 -.9 1 .4 5 ? 66.7 - 2 .• 4 -. 1 6 8 - ±- .6 1.3 2.9 7 l 66.? 3 .8 •1 8 2 60.0 -2.9 -6.8 9 3 60.4 - 2 .5 -6.4 10 4 67.2 4 .3 .4 11 5 60 .8 -2.1 -6.0 12 6 62.l -.8 -4.7 62.9 2 • •, Da ys ma.ui- Sol'y l ibra ted Mg/Kg Water l 2 2 Ex •.No. XA NTHI NE II 69.? 69.l XANTH I NE I ± -?3 - XXI I {oon) TABLE EX.NO . Days Eq uil1bra ted Sol'y Mg/Kg Gr . Mean Dev.from Dev.from Gen. Mean 'lfater XAN'rHINE III 13 l 67.? -.4 .9 14 l 70.3 2.2 3.5 15 2 6?.? -.4 .9 lo 2 62.l -6.0 -4.7 l? 3 70.6 2. ) 3.~ 18 3 71.0 2.9 4.2 19 4 68 . ? .6 1.9 20 4 -66.? 68.l - -·, .4 -.l ± 2 .o Mean of group means 6b . '/ :t 4o0 General mean 6 6. ti :r 3.0 -74 - HYDRA 1rION OF XA.NTIHNE A'l' 20°c It has been shown tha t xanthine orysta llizes out of wa ter with one molecule of water of orysta llization. (26) The follo v-1 ing experiments we re done to demonstrate whether a hydrate is formed at 25°0 . Solid xanthine wa s e quilibra te d with its aqueous solution for three or to ur days under the conditions of the solubility experiments. The solid :phase was removed by fi l tration a nd supe r f ici a l ly dried be t ween filter papers. These s amples were then submitted to drying in presence of s ulfurio a oid. solutions of kn own vap or. pressu res as i n the case of aden ine. Afte r four days t he samples we re removed , wei gh ed, and then further dried in an oven at 100°c. The res ul ts a re given in t a ble ~txIII. TABLE XXIII HYDHA 'I'I ON OF XANTHI NE Wa ter of Hydr. Wt of Desicc. VVt H2S04 soln of Wt of Superfie. Dried Xanth. Dried :Xanth. Dried Xanth. Mlvl of Hg Grams Grams Grams 1. 20 . 80 .087 .076 , (#7 50 0 t.., • •) 2 1.40 .065 . 0 49 .049 0 3. 22 . 54 .08? .066 .066 0 4. 23 . 25 . 200 . 1 75 .175 0 5. TNater .089 . 071 .066 ,5 Vap . l) ress of oven -75- It is seen from the above experiments that water is completely lost from the xanthine samples to sulfuric aoid solutions having va por pres~ures of 23.25 mm of mercury or less . In Experiment 5 the preliminary loss wa s fouod to be due to evaouation of the system for one minute. It also i ndicates that if a hydra te were formed it would have retained more than .5 moleoules of water as is indicated. From the above experi- ments it therefore appears that xanthine does not form 0 a hydrate at 25 c. -76- THE tHS SOOIA'l'TON' CONSTA .tT'r OF XANTHI N:E Arr 25oC The dissociation constant of :x:anthine was done • both by the titrimetrio and t he solubility methods. For the titration of :x:anthine 100 ml. of .0001 mola r soiution was titrated with l mn. of .ol N. NaOH. The results a:re s hown in table XXIV . There 1s an apparent drift in the values of pKa probably due to t he absorption of CO 2 from the air during the titration. It was also neces s ary to wor'J;( with extremely dilute s;::>lutions which made the titration · subject to error on this account. The average value of pKa obtained is ?.06. Because the results of the titration were not entirely satisfaotory the dissociation constan t of :x:anthine wa s also done by the soL1bili ty method. Small amounts of solid xanthine were equilibrated under the conditions of t he solubility expe ri ments with varying conoentrations of naOH solution. After three days the solub ili 1 ty of the saturated so.1.utions wa s determined. pH determinations were also ma de on eaoh solution by means of a gl ass eleotrode and a Bec klll8. n pH meter. r esults are given in ta ble 7::..01 . The The limiting factor in the preoision of these experiments was the pH determination. In the first plaoe the pH va lues tended to drift - 77 - somewhat during the determination and in the second pla ce a small change in the pH oaused a considerable chan 6 e in the pKa value. The average value for pKa of xanthine by the solubility me t hod was found to be 7.10. Combining the results obtained by the t wo methods we get pKa equal to 7.1 and the dissociation oonstant equal to ?. 9xl0 -8 . It is to be noticed that the results obtained by either method are substa ntially the same. The value for the dissociation constant of xanthine -10 c::t. -( (_' ~ :,_ is given by Wood (2]) as l. lOxlO . Such a Silk"l-11 value would be expeoted if the xanthine preparation were insufficiently pure. This same indication of insufficient purificat ion has already been referred to in connection with the solubility of xanthi ne. Furthermore judging from the ease with whioh xanthine dissolv es in NaOH solutions a s indioated by the steepness of the curve in Fi gure 6 and. by the elose s truo tura l relationship of xanthine to uric aoid one would expect the dissociation constant to be hi gher than tha t given by Wood. It is interesti ng to point out the contras t between the dissociat ion con:3 tant and the pH solubility curve of guanine and xanthineo Guanine has a small dissociation consta nt hence its solubility changes very little until - ?8- the concentration of hydroohloria acid is fairly large. Xanthine on th e other hand immediately in- creases in solubilit y even when only a s ma ll amount of NaOH has been added . From all t he evidenc e presented it app ears that our value for t he dissociation cons t a nt is more accurate than that given by 1Yood. -79- T.AB:LE XXI V ACID DISSOCIATION GONSTANT OF XA NTBI NJ1~ AT 25°c NaOH Xanthine Solution pH of Water Blank .o 5 . ~5 b . 90 .1 6 . 30 6.20 .2 6.54 6.54 .3 6. 7'7 7.00 7.l2 .4 o . 94 7.75 ? .12 .5 7.05 8.10 '7 .0 5 .6 7.19 8.50 ? .0 1 .7 7.30 8.70 .8 '7 .43 8.90 .g ?.54 9 .00 6.59 1.0 7 .66 9 .10 •• pH of pK a ral Average pJia 7 .Ob 41 • XANT'4lNE. TITRATION CURVf:5 l•O KAM,~tl'\E. o- o - o Vil\~~ °"-AN.l "--)<,. -11. CC)RlU:CTeJ> C.OlVc. :----- :I: 0 -l I .a I I I. I I )( -80- DISSOCIATION COMS TANT OF XAN'fHINE AT 25°c FHOM SOLUBILITY pH of Xanthine Solution Solubility Mg/ Kg Jater l 7 . t5 35 . 6 7.09 2 7 . 65 62.7 '7 . 20 3 7. 76 107 . 0 7 . 02 4 8.46 277 . 0 7 . 10 Average PKa 7 . 10 VARIATION bf SOLU.BILlTY 0f XANTHINt with pH 0 160 0 80 40 7 7,5 pH 8 "'F1GuRE6 URIC ACID THE SOLUBILI'll'f OF UHIC AC I D For the determination of the aruount of uric aoid in solution t he colorimetric method of Benediot and Franke (27) was adopted. This method depends ona blue color which is developed when to a uric acid solution is added araenophosphotungstia acid rea gent followed by a solution of sodium oyanideo The color is measured quantitat;ively in a colorime ter. It was establis hed by His and Paul {28} tha t the ollief d.i:f'fioulty in the determination of the solubility of uric acid is ita instability in aqueous solutions. In our experiments, therefore, urio aoid was equilibrated with its aqueous solution only for a limited period of time. That equilibrium is established in a relatively s hort time was also sl1own. by His and Paul in their work. The details of our J rooe dure are as fo .Llows. A portion of the saturated solution of uric acid was transferred from the so.Lubility tube into a weighing bottle. It wa s then treated in the same way as the stand&rd solution of urio aoid. The standard solution was pre- pared ·by dissolving 5 mg. of urio acid in 250 ml. of water. This was readily effeoted by shaking. standard was prepared each day. A new The solutions for colorimetric oom:parison wer e made up i n 50 ml. fla sks by adding, a) 10 ml. of standard uric a cid solution ..+ 5 ml. NaCN solution + 1 ml. of arsenophosphotungstio acid rea g ent. b) 5 ml . . of saturated urio a.aid solution + 5 ml. wa ter -4- 5 ml. NaCN solution + l ml. arsenophos- photungstio aoid reagent. Af ter 5 minutes these were diluted to 50 ml. with water a nd compa red in a colorimeter. For the solubility at 50°c 2 ml. of the saturated solution was diluted to 10 ml. wit h water and then treated in tp.e sa.:m.e way as oefore. The results for the solubility of uric acid at 25°0 are given in tables llVI and Y~I I. It will be seen that t he values obta ined wh en the solub_ili ty equilibrium waa approa ched from the supersaturated side are somewhat hi gher th,:i.n those obtained when t he e q u.ilibrium was approaohed f ram t h e undersa. tura t ed s i<le. This ma y be due to supersaturation or pr obably to deomnp osi tion. In the attemp t to avoid deoomp osition supersaturation oould not readily be avoided and oonversely. 'l'he solubility of urio acid at 25°c 1s taken to be .041 t 1000 grams of water. .001 grams per The molality of the saturated -83- solutio n of uric ao id. a't this temperature is taken as .0002f1-4 . The results of the det er mi nations of the solubility of uric acid a t 50°c are given in tables XXVIII and XXI:X .. The solubility is take n to be .11'7 per 100 gr ams of water. fore t aken as being -t-. 004, grams A saturnte d soL1tion is there- .oooag6 molal. There appears to be no systematic variation i n the determinations. Our res ul ts are in good a greement with those of His and Paul (28) and Gudzent (29) if a proper interpolation is made to aooount for the variation of solubility with temp erature. A summsry of the earl ier determinations of the solub ility of uric acid is given by His and Paul. Al- though it is interesting historically the failure of the early i 11ves tiga tors to reoognize t he instability of urio acid solutions malrns t heir r€sul ts of little value. -84- TABLE XXVI SOLUBILI TY OF mnc ACID AT 25°0 Equi l ibrium App roached from the Undersaturated Side Ex. No. Hours Eq uilibrated URIO ACID I Sol 1 y Dev.from Dev.from Mg/ Kg Water Gr. Mean Gen. Mean 1 8 42.7 1.9 2.0 2 22 41.6 .8 ~9 5 12 89.7 -1.l -1. 0 4 48 ~B.4 --2.4 -2 .3 5 24 41.l •3 .4 6 12 40.2 -.6 -. 5 ? 36 ~ 40.13 -~ ;~ ± 1.l URIO AC ID II 8 24 40.5 o .. o --~ 9 24 40.5 40.5 o.o .± o.o - .,::: Mean of gro up means 40.? + 1 .1 Gen era l mean 40.? :! .a TABLE XXVII 0 SOLUBILITY OF URIC ACID AT 25 C Equilibrium Approached from the Supersaturated Side Ex. No. Houra Equilibrated Mg/Kg Dev .from Gr. Mean Gen . Mean Sol'y Water URIC ACID I Dev.from l 12 39.5 -.6 -1.9 2 12 40.0 -.l -1.4 3 12 41.0 .9 -.4 4 12 40.7 .6 -.? 5 24 39 . l -1.0 -2.3 6 24 39.5 -.6 -1.9 7 24 40.5 .4 -.9 8 24 ~ -.9 40.l - -t-. 5 .4 URIC ACID II g 12 42.2 -.5 .8 10 12 42.2 -.5 .8 11 12 42.5 -.~ ') 1.1 12 12 44.5 1.8 3.1 13 24 43 .5 .8 2.1 14 24 41.6 -1.1 .2 15 24 41.6 -1.1 .2 16 24 43 . 7 42.7 .L.Q_ 1.3 :t d Mean of group means41.4 .:t • 9 General mean 41.4 --t- •. 8 -86- Ti\BLE XXVIII SOUJBILrrY OF URIC ACID AT 50°c Equilibrium Approached from the Undersatura ted Side Dev.from Gr . mean Dev.from Gen. Mean 112 -vr" -5 9 96 -11.1 - 21 3 24 100 -1 ;) -17 4 8 120 5 3 5 24 121 6 4 6 12 120 5 3 7 24 10 6 -~ - 11 8 10 116 l _1 9 24 122 7 5 10 13 126 J.1 19 11 24 129 14 12 12 ~ .l.i.J 128 13 11 13 24 1U5 -10 12 14 10 108 -7 g 15 12 121 6 11 16 12 115 -0 -c:. 17 12 118 3 l 18 12 117 2 0 19 12 11'7 2 0 20 8 112 17. -.:, -5 Sol 'y Hours Equilibrated l!ig / Kg 1 7 2 EJ!;:. No. Water UR IC ACID I ., ') 1- -87- TABLE Ex.No. :u'VIII (oon) Hours Equilibrated Sol'y Mg/Kg Wa ter Dev.from Gr. Mean Dev.from Gen . Mean 21 8 114 -1. -3 22 8 112 -3 - -5 23 8 112 -3 -? 24 8 113 115. -2 , j: 4 . -4 25 8 127 ? 10 26 8 127 7 10 27 8 124 4 7 28 8 125 3 b 29 8 124 4 ? 30 12 11~ -1 2 31 12 124 4 ? 32 12 124 4 ? 33 12 128 8 11 34 l2 128 8 11 35 8 10'7 -13 -10 36 8 106 -J.4 -11 37 8 111 -~ -6 38 8 111 -~ -6 39 8 -110 120 -:t-10 Mean of group means 117 + 6 General mean 117 --t- muc ACID II 4 3 -? -88- rAJ:lbE 1 XXII SOLUB I LI TY OF URIC ACID AT 50°c Equili brium Approache d from the Supersa tura te d S ide Ex. No. Hours 11:q ui- Gr. lvlea n Dev.from Gen. Mean 12 0 3 2 6 118 l 0 3 6 l-2 2 5 4 4 6 l2 5 8 ? 5 24 109 -8 -9 6 24 113 -4 -5 7 24 115 -2 -3 8 24 115 ll'i ~ -2 +4 -3 9 6 116 -3 -2 10 6 1 14 -5 -4 11 6 114 -5 -4 12 6 115 -4 -3 13 24 123 4 5 14 24 110 -9 -8 15 24 130 11 12 16 24 -119 - 10 -t 5 Mean of group means 118 -+ ? General mean .± 3 11brated Sol'y Mg/ Kg Water 1 6 2 'ORIO ACID L Dev.fr om UR IC ACID II 128 118 9 THE DISS OCIATION CO:NS 'f.' NT OF UilIC ACID AT 26°0 The dissociation constant of uric a oid was deterro.ined by electrometrio titration. 100 znl. of .000238 molar uric aoid solution was titrated with 1.7 ml. of .0143 N . Na.OH . The results of the titration are shown in ta'b le X:X:X . It will be seen that in the middle portion of the titration the pKa values are in good agreement with one another . Urio aoid is a fairly strong acid and can therefore be titrated satisfactorily in spite of its slight solubility~ The average value of pKa obtained from the titra. tion is 5.76. The dissociation constant is there- fore taken as l.74xl0 -6- • measurements at 18°c. From careful oonduotivity His and Paul (28) found the dissooiation constant to be l.5lxl0 ... 5 , a result which is in substan tial agreement with our value. Musgrave (30) give 8xl0 stant of urio aoid, -o Young and tor the dissociation oon- Their value was derived graphically from titrinetric datai TABLE XXX 0 DISS OCIATI ON CONS TANT OF URIC ACID A'r 25 C NaOH pH of Urio Acid Solution pH of Water Blank 4.96 6.80 .... .l 5 .06 7.51 6.02 .2 5.18 8.40 5.94 .3 5 .28 8 .7? 5 .8'r .4 5.38 9.00 5.8~ .5 5 .46 9.12 5.77 5.56 9.22 .? 5.64 9 .29 5 .86 .8 5.72 9 .37 5 . 75 .9 5.82 9.44 5 . 74 l.O 5.94 9.48 5 .76 1.1 6.01 1.2 6 .12 1.3 6.28 1.4 6.41 1.5 6.56 5 . 50 1.6 6.81 5.38 1.7 7.35 e o • o 5 . 74 Average pKa 5.76 l-6 I ./ URICACID . TITRATION CURVES / U~IC ACIJ> 0- 0 - 0 WAT"" &LANK • -•--"' -. ,.o I I .s ' I . .s J I A- J( ,. o&,.....-_ _......__ _ _....___ _..._,L__ _ _~ - - ~ : - - - _ _ _ , , 4 5 GpH7 'f'tG.UR.E..1 9 - 91- ALLANTOI N THE SOLUB I LI TY OF ALLANTO I N 1he solubiLity of allantoin was determined by the gr§ ~im_etrio method . The result of t he s)l ub111ty de- termi na ti ons a t 25°0 are given in tables A"XXI and XXXII and at 50°c in t a bles XXXIII and X::Z:XIV. It appears t ha t .All an toi n II has a s ma ller solubi l ity. enoy of a llantoin to s upersa turate The tend- is also notioeable. Whi l e there are some large disorepanoies among the indivi dua l r esults tbe solub ility of allantoin is sufficien tly large to make these deviations relatively s mal l. The so l ubility of allantoin at 25°0 is taken to be 5.92 ::t • J is .0373. . 1'he m:.:,la li ty •) f the saturated solution The solubility of allantoin at 50°c was found to be 18 .L:t•3 grams per 1000 grams of water. The mola lity of the saturated sol~tion was taken to be .114 • According to Marti gnon (31) the solubility ot allantoiu in 1iv ater is 7.7 grams per 1000 grams of water at ordinary temp erature. Schulze and Barbieri (32) give 5. 4 gr ams per 1000 grams of water at 22°c a result which is on good agreement with ours . .A lla.ntoin is also re 9orted a s e asily soluble i n hot water. -92- TABLE l.-XXI SO LUBI LITY OF Al.LANTO I N AT 25°0 Equilibrium App roached from the Undersaturated Side Ex. No. Equilibrated Gin/Kg Water Gr. Mean Dev.from Dev.from Gen.Mean Sol'y Days ALLANTOIN I l 2 5 . 99 -.01 .13 2 3 6.00 .oo .14 3 4 o.og .09 . 23 4 2 5. 89 -.11 .03 5 4 5 . 98 -.02 .12 6 5 5.94 -.06 .08 7 6 - 6.08 .os .22 6.00 ± .05 ALLANTOIN II 8 2 5.73 .01 -.1:3 9 3 5.72 .oo -.14 10 4 5.'75 .0 .,,. -.11 11 2 5. 67 - .O \/ - •i9 12 5 5.70 -.02 -. 1 6 13 6 5.68 -.04 -. 18 14 7 5.76 5 . 72 .04 + .03 -. 10 of group means 5.86 + .06 General mean 5.86 -t- .08 lVlean ' TABLE XXX II SOLUBILITY OF A.LLANTOIN AT 25°0 E~uilibrium Approached from the Supersa t urated Side Ex. No. Days Eq_ ui- Librated Sol'y Gm/Kg Dev.from Gr. Mean Dev.from Gen. Mean Water .ALLANTOI N I l 3 6.17 .18 .19 2 4 6. 0 5 .06 .07 5 5 5.d9 -.10 -.09 4 6 6.00 .01 .02 5 6 5.90 -.09 -.08 6 ? 5.93 -.06 -.05 ? 8 5. ~n -.02 -.0 1 8 9 6.03 .04 .05 9 11 - 5.98 5.99 -.0 1 + .06 .oo 10 3 6. 34 .37 . 36 11 4 6.09 . 12 . 11 12 5 5.96 -.01 -.02 13 6 5.92 -.05 -.06 14 7 6.16 .19 .18 1!5 8 5.g 5 -.02 -.03 16 9 5.84 -.13 -.14 17 10 5.7? -.20 - .21 18 12 hZ! --.26·~· -.27 ALLANTOI N II ± . 1;3 Mean of group mea ns 5.97 5 .98 General mean 5.98 :t .07 :t . 14 TABLE XXXIII 0 SOLUBI Lirr y OF ALLANT OI N AT 50 0 Equi.Librium Approached from the Undersaturated Side Ex . No . Days Equilibrated Sol'y Gm/Kg Gr. Mean Dev.from Gen. Mean Dev.from Va ter ALLA NTOIN I l 1 17.81 -.04 -.07 2 2 l?. 71 -.14 -.17 3 3 17.88 .03 .oo 4 4 17.92 .07 .04 5 5 17.93 .,08 .05 6 6 J.'7. f35 .oo -.03 7 7 17.86 17.85 -06 -.02 .01 ::t. • ALLUJ'rO I N II 8 2 17.74 -.15 -.14 9 3 17.76 -.13 ,-.12 10 4 18.08 .19 .20 11 5 18.22 .33 • 12 e 17.67 17 . 89 • -.22 ± .26 Mean of group means l? . 88 -+ . 2'J General mean l?.88 :±. .·, 0 Iv 3 !.) -.21 - 95- TABLE mrv SOLUBILITY OF ALLANTOIN AT 50°0 Equilibrium Approaohed from the Supersaturated Side Ex . No. Days Equilibrated So.l'y Glll/Kg Water ALLANTOIN I Dev.from Gr. r. e ean l l 17.?3 -.26 2 2 18.12 .13 3 3 18 .09 . 10 4 4 18 .11 .l •) 5 5 1 7.99 6 6 17.90 17.99 7 l 18 .85 8 2 17.68 9 3 18.51 10 4 lf3.53 11 5 18.47 12 6 18.70 .:, Dev.from Gen. Mean -. 10 -.11 ALLA NTOIN II 18.46 Mean of group means Ge neral mean 18.23 -. 5 4 - . 2 ··-4 ± .30 .4 b THE DISf:lOO I A'l1I ON COWS T l\'NT OF ALLANTOJili AT 25°0 To determine the dissociation consta n t of alla ntoin 10 ml. of .038 molar solution was titrated w1 th an equal volume of .038 N. Nao~ . The results of the titrat ion ,9. re g iven i n t abl e :a..,T I . The titra tion curve is shown in Fi g ure 8 . It will be seen that the pKa, values are quite eonsiste nt so tha t an a verage value sho uld give a good • dissociation constant. The a ve rage pKa value wa s found to be 8.6(/ whi ch g ives tor t he dis s ociation const a nt of Hllanto in at 25°c Ka .:::: 2 & 1 4 d O- 9 • Vl ood {~·l .i ) determined t he d issociation constant of allantoin by followin g the rate of oa t a lysjj.s of the saponification of methyl a ceta te by allantoin hydrochloride. His value is Ka ~ 9 l .17xl0 - . The ti trimet- rio method is to be considered to be mor e rel iable. Wood 's work was done at a time when t he mea. suremen t of H had not yet been developed to a hi gh degree. His determination gives a result of the s ame order ot magnitude a s ours. TABLE x.xxv DISSOCIATI ON CONS TANT 01r J.\. l.LA N'l'O I N AT 25°c NaOH pH of Allan·toin pH of Water Blank pKa ml Solution .o 6.21 5 .16 . ... . 25 ?.14 10.66 8.73 .50 7.46 10.99 8.62 .?5 7.63 11.14 a.?;) 1.00 7.73 11.24 8 .67 1.50 ? • 9;5 11.39 8.68 2.00 8.06 11.46 8.66 ::! • 50 8.16 • 0 8.64 3 . 00 8 .28 ll .56 8.65 4.00 8.62 11.64 8.80 5.00 8 . 62 11.66 8.62 6.00 8.81 11.72 $.64 7.00 9 . 00 11 . 73 8.63 8 .00 g .21 11.75 8.62 9 . 00 9.52 ll .'76 .. . . 10.00 10.00 ll.'78 •••• • • • pKa '9.6'/ Ka. 2. l 41:10 -9 10 ::c ¾ ALLANTOIN TITRATrON CURVES 41.L ANTOIN •-o - o / WA'tfR BLANK. -"- ~ - ~ CORRE.C:T~J:, CURVE - - - - / / I I i: I 1/ 8 V 1/ 6 4 · 2 0 L..~:::::::::1:::__-==='=~ ·e 7 / I =:::±=======t====---_L__ _____: g (0 ALLOXAN THE SOLUBILITY OF ALLOY~N There are no quantitative data in the literature on the solubility of alloxan. The solubility of allo:x:an was determined gravimetrioally. There are several points of departure from the general method whioh should be noted. Since alloxan is very soluble only small portions of the saturated solution were required for a single determination. .5 grams of the solution wa s used. About If the samples were larger some diffioulty was encountered in the drying. Drying was done in vacuo over sulfurio acid at room temperature in order to prevent decomposition. It required about a week for the samples to dry to constant weight . Unde r these conditions alloxan contains one molecule of water of constitution which is sometimes referred to as water of crystallization. In order to obtain reliable and cons is t ent results one must be aware of the difficulties enoountered d;ue to the instability of a lloxan. It decomposes to urea, oxalic acid, alloxantin and carbon dioxide even when dry as was shown by Gortner (33). When in a closed system several explosions have bean reported (34), (35), (36). In aqueous solutions this• decomposition is accelerated considerably which makes the determination of any physical constants of alloxan somewhat difficult. Freshly crystallized alloxan was used in all the exp~riments and the ti me of oonta.ot with water was also deoreased to a minimum. The results cf the solubility determinations of alloxan are given tn tables XXXVI, XXXVII, XXXVIII and XX.XIX. At 25°C the solubility wa s found to be 416 -;:. 2 grams per 1000 grams of water,, corresponding to 2. 60 molal solution at saturation. At 50°c the solubility was found to be 719 .;t:. 2 grams per 1000 graxns of water ,oorresponding to 4.4 9 molal solution at saturation. Because of the hi gh solubility it is rather easy to obtain results with considerable precision. Some unoerta.inty is introduced on account of the instability of alloxan al thou gh this has been reduoecl by work- ing under favorable oonditious .. Note- the molecular weight of the stable modification of alloxan at 2)°C is taken as 160.o~·,. This corresponds to a formula C4H4~,Nt• -100- TABLE XXXVI SOLTIBILITY OF ALLOXAN AT 25°c Equi l i br i um Approached from the Unders a tura ted Side Ex. No. 1 Hours Equi- Solubility Dev. from librated Gm/ Kg Wa ter the Mean 24 417. 6 1.0 41~ . o -.6 2 3 416.8 " 4 5 " Mean 'rABLE 416.2 -.4 416 . 5 -.1 416. 6 + .b XXJCVII SOLUBILITY OF ALLOXAN AT 25°0 Equilibrium .Approached from the Superse.tura·ted Side Hours Eq_ui- Solub ility libra ted Gm/ Kg l:Ya te r Dev. from t he Mean l 24 418.5 2.4 2 ft 415.5 -.6 3 " 416.6 .5 4 ?t 415 .3 -.8 5 ti 414 .6 -1.5 Ex. No. Mean 416 . l I 1 .•4 -lOk TABLE XXXVIII SOLUBILITY OF ALLOXAN AT 50°c Equilibrium Approached from the Undersatura.ted Side Ex . No. Hours Equilibrated Solubility Gm/Kg Water Dev. from the Mean 1 3 717.1 -2.0 2 It 719 .5 .4 3 " 720.7 1.6 719.4 .3 1718.9 -.2 4 5 ,, It ?U~ . l :i:_ 1. c Mean TABLE XXXIX SOLUBILITY OF ALLOXAN AT 50°0 Equilibrium Approached from the Supersaturated Side Ex. No. Hours Equilibrated Gm/Kg Water Solubility Dev. from the Mean l 3 718.5 -.3 2 " 718.4 -.4 3 it 717 . 9 -.9 718 .9 .1 720.3 1.5 4 5 ,, ,, Mean 718.8 - -+. 8 -102- TH'E uISSOCIATION CONSTAN'.i"S OF ALLOXAN AT 25°0 The dissooiation constants of alloxan were determined by eleotroraetrio titration. . 10 In.lo• of .0937 molar alloxa.n solution was titrated with 10 ml. of .0989/o'N. Ne.OH for the first dissociation a onstant an d an additional 10 ml .. of Na.OH for the second dissociat ion cons tant .. During the first part of the titration the pH values were observed to drift especially if a suff icient length of ti me were allowed. In order to obviate this difficulty the ini •tial pH readings were taken.. During the seeond part of the titration of the firs t dissociation constant Ocbd pH readings were steady potentials were obtained.. also obtain ed throughout the titration of the seoond constant. A considerable uncertaint y ha.s ·therefore been introdueed into t .h e value of the first dissociation constant due to the instability of a lloxan. An average of the values obtained in the middle portion of the titration was ta~n. pK was found to be 6.'L() giving a -7 for the dissociation constant ~.~9xlO • The results are shown in table XL. The seoond dissooia tion conste.nt is much more satisfactory as may b e seen from the data in table XLI. An average of the pKa values gives 8.S6.lfhe dissociation oonstan t comes out to be 2 .l9xl0 -9 • - 1.03- The a_issooia. tion o onstant of alloxan was determined -...ir-:~ by Wood (,'f~ and by Trubsbach (:37) who give 2 . 32:xl0 and 4.llxlO -5 • .. 7 Both investigators used the oonduotivity method and both found that the conductivity changed during the determination. The value given by Wood is of the same order of magnitude as our first dissociation oonstant. By the conductivity method the seoond dis- sociation constant was not detected. The titration curve of alloxan is given in Figure 9. 20 ALLOXAN TITRATION CUR.VE.S ALLO" A "i - " - " , WATE"..!,L-4N~ ,<-"--~ CORRe.C'Tf.JJ CURVI?. - - - - • . - 12. 8 OL...,..__===::::::::::__L-_ 3 s _.,~--='========:=..:.~ pH 9 -104- • TABLE XL THE FI:i~ST DISSOCIATION OONS 1rANT OF ALLOXAN AT 25°0 NaOH pH of Alloxan ~olution pH of Water Blank o.o 3.50 7.36 •••• .5 5.90 11.50 7.16 1.0 6.17 11.70 7 .10 1.5 6.25 ••••• 6.98 2.0 6.32 11.85 6.87 * 2.5 6.40 ••• •• 6.86 3.0 6.41 11.90 6 .,75 3.5 8.48 ••••• 6.71 4.0 6.57 ••• 6.71 4.5 6.64 • 6.68 5.0 6.69 6.63 5.5 6.74 6.59 6.0 6.83 6.59 6.5 6.93 6.58* 7.0 6.99 6.54 7.5 7.08 6.50 8 .0 7 .17 6.44 8.5 7.31 6.37 9.0 7.45 6.20 9.5 7.60 •••• 10.0 7.75 •••• ml PK1 6. ·10 -105- TABLE XLI THE SECOND DISSOCIATION OONSTANT OF ALLOXAN AT 25°0 NaOH pH of 1\ llo:x:an Solution pK2 10.0 7.75 9.21 10.5 7.90 8.82 11.0 · 8.0l 8.?5 ~ 11.5 8.13 8.70 12.0 8.23 8.66 · 12. 5 · 8.32 8.65 13.0 8.42 8.63 13.5 · 8.51 8.65 14.0 8.60 8.65 1-4.5 8.69 8.64 15.0 6.79 8.65 15.5 8.89 8.65 16.0 8.99 8.65 15.5 9-10 8.65 17.0 9.23 8.64 17.5 9.41 a.ea 18.0 · 9.&3 8.64 18.5 9.93 .... 19.0 10.51 •••• ml pK2 8.66 -106- VAPOH VitESSUHE OF ALLOXA1J- ALL OXAN THIHYDH.ATE SYSTEM AT 25°C What is usually referred to as allo:x:an monohydrate is the stable modification of alloxan at 25°0 and the water of crystallization is aotually water of oonstitution. It will be referred to simply as alloxan in con- formi ty with the terminology of Stiehler and Huffman {l) , .( 2) ·• On this basis the tetrahydra te becomes a trihydrate. The vapor pressure of alloxan-alloxan trihydrate has been determined by the method previously described. Large alloxan trihydrate crystals are easily prepared by crystallizinB alloxan from water. f3 uoh crystals were removed from the mother liquor and superficially dried between filte r papers. Because the cryst als are large the water of wetting can be almost comp letely removed in this manner. Crystals of alloxan trihydrate were submitted to desiooation in presenoe of solutions of sulfurio acid of known vapor pressure. The loss in weight of the alloxan trihydrate orystal was determined after four days. The results are given in table XLII. The dehydration is slow so that the loss of water was not oom:plete in all oases. The loss, however , is s ufficien t to in1icate wheth er the vapor :9ressure of the hydrate system is greater or less than that of the -107- sulfuric acid solution. The vapor pressure of alloxan- alloxan trihydrate system is taken as the average of experiments 4 and 5, i.e., 21.l + .3 mm of mercury. Alloxan was als o shown to hydrate readily when the anhydrous form was placed in :presence of water 0 vapor at 25 c. TABLE XLII VAPOR PR ESSURE OF ALLOXAN HYDRATE SYSTEM AT 25°0 No. Wt of Allox. Hydr . Xtal After Mol. of H2S04 Sol. Desiccat. Rema ining mm gms . gms . Vap.Press. Wt Water Is•· . I-/ l 10.9 .682 •-459 2 17.8 . 372 .274 .o 3 19.4 .186 .149 • '7 4 20.8 .075 .Ool .8 5 21 . 4 .0?5 .075 3.0 6 22 .4 .208 .210 3.0 .9 -108- ACTIVI 'rY COEFFicnm1rs OF ALLOXAN' AT 25°c Because alloxan is unsta ble in a queous solutions a suit abl e method was developed for dete r mining the activities under such conditions. The method depends on the principle of isothermal distillation. The pro- cedure wa s as follows. The apparatus con sisted of a small desiccator with water at the bottom so that the atmosphere inside was satur ated with water vapor. Into this desiccator were pla ced s mal l glass cups of ap proximately the same dimensions and½ ml. capacity. They were placed on a s uita ble coppe r bloc k with their places numbered. Into t wo such vessels were placed .25 ml. of standard NaCl solutions, the oonoentrations of wh ich differed by about 10%. Into t wo other cups were placed .25 ml. of alloxan solution of suoh a oonoentration that it s vapor pressure wa s between t hose of the two standards. It was possiole to select th.is conoentration by as s uming eqlLl- the behavior of alloxan to be the same as that ofAmolar urea for wh ic h the vap or pressures are known (38). Sometimes prelimi na ry runs had to be made in order to determine the proper oonoentrations to be used in the actual expe·ri ment. -109- The glass cups we re weighed with their respective solutions. Care was taken to prevent evapora tion during the manipulations by covering the vessels with a glass cover. The vessels were wei ghed uncovered. The vessels were then pla oed on t he copper bloak and put in the desiccator. bator. The desiooator was placed in a 25°0 incu- The length of time required for an appreciable amount of distillat ion to take place was about 12 to 15 hours. For the more ooncentra ted solutions the time could be shortened to about 8 or 10 hours. The temp erature in the incubator remained constant to within .1° or less. After a suitable period of time the vessels on the cop per block were removed from the desiccator, covered with a g lass cover and finally weighed after about 5 minutes. This time was requi r ed for the water film on the outside of the vessels to evaporate in order to increase the aoouraoy of t).}e weighings. An example will illustrate the details more clearly, TABLE XLIII VAPO R PHESSURE DET ERMI NATION Solution rJJolali ty Initial Wt. Final wt. Gain grams grams mg. NaCl 1.130 .7921 .7964 4.3 NaCl 1.300 .7448 .7516 8.8 Alloxan 2.500 .73?? .7432 5.5 Al loxan :?..!'500 .7980 .8032 5.2 -110- The initial and final weights refer to the total weight of the vessel plus the solution. 81:noe it is the gain in weight that is required the absolute amount of the solution need not be known. The volumes of the solutions were measured with a pipette and were .25 ml. in all oases. The average gain of the allo~an solution from the above data is 6.35 mg . The agreement between the two duplioates was usually better than in this example. It will be seen that 2.500 molal alloxan solution is isotonic or isopiestic with a NaGl solution which is between 1.130 and 1.300 molal and by interpolation this oomes out to be 1.194 molal. On the average the results obtained agreed within 1%. In this manner the isotonic solutions corresponding to four different aoncentrations of alloxan were determined. Eaoh result was obtained as the average of several determinations. data obtained.. Table XLIV oontains the -111- TABLE XLIV !SOTO IC NaOl AND AU.OXAN SOLUTIONS Molal1ty of Molali ty of Isotonic faCl Value or 1668 • 700 .3803 3.27 23.460 1 . i50 .62? 7 3.27 23.269 .9294 3.29 23.034 3.·32 ' 22.827 .Allo:mn Soln. Vapor .Press • Soln . I ' I 1.8575 l.'1..btn 2.500 I w 1 .1960 The oaleulation of vap i:, r pressu:res was made by the use of the formula l00R :.. l00<:120-Pl Mpo given in the I . C. T. where P 0 is the vapor pressure of pure water, p ,is the vapor pressure of NaCl solution and ~ i.., M is the mole.11 ty. B 1.s defined by the equation and is the fraotional vapor pressure lowering per mole ot solute. The values of R for NaCl were obtained •f~om I.C.T. A graph ot vapor pressure against molality was made for alloxan using the values trom the table above. is shown in Figure 10. This Aotivity ooettioients were oal- oulated for several oonvenient rounded oft molalities aooording to the method of l..ewis and ~andal.l {39) using the h/m funotion. In table XLV the results of these oaloulations are shown. VAPOR PRe.~~URE..5 of ALLO"AN SOLUTIONS a-t 2 s·c 2.3,00 I 23,80 , - Z ♦,oo'-----------'------"---------11-------...L----~--=--- o I M o '-ALITY 2 of AL'LOXA ~ SOLUTIO N 2,-S,.,. Re «>IL.. ~ »-n1,(. ~ ~ ~~ = • 8 I ::t. • .,_ ct > ·J'l .... ( ~ s J-1.( 4'?:...) ~ ~ ~ ~ O r . ~-•~ <-<e.e -C....· - 112- TABLE XLV ACT I VITY COEFFICI ENT.S OF ALi.OLUf Molali ty of Alloxan Soln. Vapor Pres s . Mm of Hg b/m .50 23.54 5 . 0 092 -~~84 1.00 23 . 340 .0138 ,9953 1.50 23.140 .0274 .9880 2.00 22.967 .0315 .9809 2.50 22.810 .0393 .9'l0-5 2.60 22.780 .0403 . -9 -682· Activity coefficient In order that the method of calculation be made clear all steps of the calculation of the activity coeffioient for satu1·a ted a lloxan so lu tion are given. The vapor p ressure of 2.60 molal alloxan solution is obtained directly from t h e graph in F igure 10. value is 22.780 mm. Hg . This The vapor pressure l owering is given by the formula p 0 - p p which ~mounts to .04108. 0 The vap or p ressure of wa ter is a measure of the fu gaoity or activity of t h e solvent a 1 • The value of ln a 1 is next obtained by means of the fo.rmula - p 0 -p -113 - by substituting .04108 for the fra.otional vapor pressure lowering. ln a 1 is thus -.04193. The h funo tion is rela ted to the ao •ti vi ty of the solvent by h ... 55.51 11/\Cll + l m where mis the molality and is 2 .60 in this case. This gives a value of .1049 for hand .0403 for h/m. The integral of h/m from Oto mis obtained by graphic integration as shown in Figure 11. i ntegration gives -.05793. Suoh graphic This value together with the va lue for his substituted in the e quation ";,,{ ln a 2 /m ::::. r -h- J h/m dm I) , 't '· !~ to g ive the aoti vi ty coefficient of alloxan of .-9&82 when evaluated. It is to be noted that the h/m function is very sensitive so that some oare is necessary in drawing the va por :p re s sure curve in Figure 10 in order to avoid tremendous distortion in the h/m values. Other values were oaloulated in exactly the same manner . bRAPHIC. !NTf.~ RA"TION b\ h/m. l=°UNC.TC O fi . • \ f\l\C> LALi TY 2. -114- l.EUCINE THE S OLUBILITY OF L..'!lUCINE I'E has been shown by Muller (3} that leucine prepared f.rom natural sources is invariably oon~inated with methionine. Fox (1~) has been able to prepare methionine free leucine. Our determinations on the solubility of leucine were done on methionine tree preparations. Both the op tical isomers -..vere use.d. The solubility of d- and 1- leuoine at 25°0 was found to be 22.0 ~.2 grams per 1000 grams of water. The results are given in t8bles XLVI and XLVII . The p recision of the results is y_Ui te satisfactory. There a ppear to be no systema tic variations. The molality of saturated leucine solution at this temperature is .16'j . 0 The solubility of d- and 1- leuoine at 60 C was found to be 26. 6 ::t .3 grams per 1000 grams of water the molali ty of the saturated solution at this temperature being . 2 02,. The results obta ined by a pproaching the e q uilibrium fro m t he s..ipersaturataii are somewhat higher but the. discrepancy is not large. The results given by Dalton and Schmidt {4} on the solubility of 1- leuaine are about 10% higher than our values. One would expect a hi gher value for the -115- solubility if the prepara tions were contaminated with methionine . The work of other investiga tors is subjeot to t he s ame criticism. The pH of a s a tura ted solution of leuoine at 25°C was found to be 6.5. -116- TABLE XLVI SOLUBILITY OF I.EUGINE AT 25°0 Equilibrium Approached from the Undersatura ted Side Days Eq_u ilibra ted Sol•y Gm/Kg Water Dev .from Gr. Mean Dev.from Gen.Mean l 2 21 . 86 -.17 -.17 2 3 22.14 .11 .11 3 5 ,~2 . 02 -.01 -.01 4 7 22 . 06 .03 .03 5 i3 22 . 06 22.03 + .03 .09 .03 6 ,. C') 21 .88 -.14 -.15 7 4 22 .12 .10 .09 8 5 22.06 .04 .03 9 8 2~.02 .oo -.(Jl 10 9 22 . 03 .01 .oo 22.02 :t. • 08 Mean of group means 22 .00 :t .11 Genera l mean 22.03 :! .oo Ex . No. d-.Leuoine 1-Leuoine -117- TABLE XLVII SOLUB I LITY OF LEUCI NE AT 25°c Equilib r i um Approaohed from t he Supers atura ted S ide Ex. No. Dev. from Gr . Mean Days Equi- S ol'y li bra. ted Gm/Kg Water 1 ~ 22 . l? . 25 .13 2 19 21 . 88 -.04 - • .l.6 3 21 21~73 - • .L9 - - ~i 4 22 L l ._59 -. 23 - ~~-S5 5 23 2B.14 . 22 .10 21 .92 :± .20 Dev.from Gen. Mean d- Leucine 1- Leuoine 6 2 22 .17 .01 . 13 7 19 22 . 21 .05 .l? 8 21 21 . 98 -. 18 -.06 9 22 22 . 10 -.06 .06 10 23 22.35 .19 .31 22 .16 .:t .12 Mean of group mean 22 .04 j- . 20 General mean 22 .04 ± .14 - 118- TABLE XLVIII 0 SOLUBILITY OF L:EUCINE AT 5o o Equi librium Approached from the Undersaturated Side Ex. No. Dev.from Dev. from librated Sol'y Gm/ Kg Water Gr1-) Mean Gen. Mean l l 26. 61 -. 17 .02 ,:.., ') 2 2 6.64 -. 14 .05 3 3 27.14 • .) 3 . 55 4 4 26 .04 -. '74 .... 55 5 5 27.11 . .j:j .52 6 6 27.14 26.78 .)6 .55 .:t . 3 '4 Da y s J.i:qui- d-Leuci ne 1-Leucine 7 l 26.48 .08 -.11 8 2 2 6. ~SQ - ~01 - . ~o 9 3 26.19 -, 21 -.40 10 4 2 6. 42 , . 02 -. 1 7 11 5 20 . 33 -,Q_'l. -.26 12 6 26.58 . 18 -.01 26 . 40 -f e l 'l Mean of gr oup means 26 ,,'•.,9 .± . 3b Gen eral mean 26.59 r . l:'.c -119- TABLE XLIX 0 8 0 .LlliH.LITY OF LEUCINE .AT 50 0 Equil ibri um Approac hed from t11e Supersaturated Side Days Equilibra ted Sol'y Dev.from Gr o Mean Dev .from Gen.Mean 1 2 26.40 -.33 -.27 2 3 26.64 -.09 - ·.03 3 4 26 • '. Jb .03 ·. 09 4 5 26.88 .15 . 21 5 6 26.82 .09 .15 6 7 26.88 26.73 .15 .21 -±. .15 Ex.. No. d-Leucine Gm/ Kg Water 1-Leuoine ? 2 26.59 -.02 -.08 8 3 2 6 .68 .O? -.0 1 9 4 26.77 . 16 .10 10 5 26 .63 .O') -.04 11 6 26 . 79 . 18 12 ? 26 .22 26.61 - .39 :t .l? Mean of group means 26.72 :t, • L j General mean 26 . 6 I ± .L~ ~ - • 1 ~, jj -. 4 5 - 120- 0 T HE DI SSOCIAT I ON CONST,\ NTS OF 1-l,EUCINE AT 25 C The dissociation constants were determined on methionine free 1-leuoine. 10 ml . porti ous pf .0972 molar solution were titrat ed with 10 ml . of .0989 N. NaOH and HCl res pec tively . The results have been calculated on the basis of the olassioal theory and are given in tables Land LI. A titration curve is s hown in Figure 12. The avera ge PKa value was found to be 9.70, the -10 d issoaiati an constant therefore being 2 .0xl O • This value agrees fairly clos ely w1·th the values found by Sano (40. ) , Harris (41 ~ , and Winklebleo h (42). Wood (43) gives a value much sma ller than this. The average value of pKb obtained i s 12 . 4 g iYing for t l1e dissociation constant Kb 3 .98xl0 -13 . This result is muoh smaller than t hat given by the above mentioned investi ga tors. The determina t:i.on of dis- sooiations constan ts in this range is subject to several prac tical diff ic ,11 ti es. However the precision of the res ults in our determ ina tio.ns appea rs to be goo ri a n d the d isorepanoy may be due to the methionine contamination. Methion ine has a basio dissociation constant in this range and would be a likely source of error. The zwitter ion consta nts for optically active leucine from the a bove res ults are K8 ::: 2. 5lxl0 -2 • and Kb = 5 . 0 lxlO -5 . - l22 - TAB.LE L BAS IC DISSOCIATION COJTST.AN'11 OF 1-l.EUCI NE AT 25°0 HCl ptt of .i.,eucine Solution pH of Water Blank .o 5.54 6.80 ..... .5 3 . 04 2 . 53 12 . 3 1 .0 2 .68 ~ . 15 12 . 4 1.5 2.55 2.00 12 .3 2 .0 2.38 1.87 12 .4 2.5 2.27 1.80 12 . 4 3 .0 2.18 1.72 12 .4 3.5 2 .10 1 .65 12 . 4 4.0 2.05 1.60 12 . 4 4.5 2 .00 1.58 12 . 3 5.0 lo9 3 1.58 12.4 5.5 1 .88 •••• 6 .0 1.83 1.58 6.5 1 . 78 •••• ? .O l.'7 3 1.54 12 .f) '7 . 5 1.68 .. . . 12 .5 8.0 ..1.. t).) .... <?.: 1 . 45 12 .0 8 .5 l.oO •••• 9 .0 1. 57 1. 43 12 . 6 9 .5 1.54 •••• 12 .6 10 .0 1. 52 l.38 •••• , pKb 1 !:'.. 4 pKb 12.4 12 . 4 12. 4 12.6 - 123- TAB.LE LI ACID DISS OCIATION CONSTANT OF l-U7JCINE AT 25°0 Ne.OH pH of .Leuoine Solution pH of Wa t er Blank pKa .o 6.32 7.50 •••• .5 6.46 11.50 9.72* 1.0 8.77 11.70 9.72 1.5 8.96 ••••• 9.71 2.0 9 .12 ll.85 9.?l 2.5 9.23 ••••• 9.70 3.0 ~.35 11.90 9.71 3.5 9.46 9.?2 4.0 9 . 53 9.72 4.5 9.65 liJ.71 5.0 9.?0 ij.68 5.5 9.77 9.66 6.0 9.88 ~.68 6. 5 9.~n 9.66 7.0 10.07 9.67 'l. 5 10.19 9.75* 8.0 10.27 9.63 8.5 10.42 9.62 9.0 10.58 9.55 9.5 10.78 9.30 10.5 11.03 •••• pKa 9.70 ,,o , LE.UCfNE TITRAT ION CURVES LWCINE o - o - o WP.Te.( 8&..ANi< (- "--111. COftft&!CTE.1) CUfCVEt - - - - - s 0 :c. - ~ 2. I 0 :c. 2 0 ~ z 6 pH -124- • LII TABLE SUMMARY OF SOLUBILITY DATA Compound 0 Solubility at 25 C Gm/ Kg Adenine 1.13 :t Solubility at 50°c mater Gm/Kg .04 3.4 ±. .2 2.01 + .06 Water Hyp oxanthine .73 :t .01 Guanine ,0034 -+ .000.1 .0131 :t . 0004 Xan thine .01_64 .± .0007 .06'7 + .ooA- Uric .Aoid .041 ::t' .001 .117 + . 00.4 .± •1 + .2 Allan to in Alloxan 5.9 416. ::t: ,:! ~ 18.l 719. J.. 2, d- and 1Leuoine 22.0 :;f:: .2 26.6 --1. •j -12 5- TABLE LIII SUMMAHY OF DI SSOCIATION OONSTAN''L'S AT 25° Compound Kb Ka .A denine l.26xl0 Hypoxanthine l.07xl0 Guanine •••••• Xanth i n e '7 . g xlO -10 -9 -10 -13 2 . 0 :xlO - 11 1 . 3 xlO -8 ••••••• -6 Uric Aoid l.74xl0 Allan to in 2 . 14x10 Alloxan ~.o xl0- 7 .. . . . .. -9 2 .19 xlO 1- Leuoine 1 . 3 ~ '!10 0 -9 2 .OOxlO-lO • • • • • • ••••••• 3. 98xlo- 13 -126- TABLE LIV SUMMARY OF OTHER DATA pH of saturated solution of adenine at 25°c 6.6 pH of saturated solution of hypoxanthine at 25°c 6.4 pH of saturated solution of d- and 1- leuoine at 25 0 C 6.5 Vapor pressure of alloxan - alloxan trihydrate system 2l . l ± .3 mm. of mercury Vapor press ure of saturated solution alloxan 22.78 mm. of mercury The activity ooeffioient of alloxan in saturated solution at 25°C .9682 -12?- DERIVED DATA The following general considerations of the methods used in the c alculation of the free energies of formation of the ionic and the non ionic s p ecies and other pertinent data are g iven here f or reference o One example o f ea ch oaloulation is given in detail in order toillus tra te the appli cat ion of t he methods. The heats of solution are calcula ted for t he temp erature range 2 5 50°c. All the other data are calculated for the tempera - ture 298.lK. THE HEAT OF SOLUTI ON If the solub ilities at two temperatur es a r e known the heats of solution may be readily derived by the ap p lication of the van't Hof f equation, d.lnK 7iT = whe re T is the absolute tempera t u.r e, R is the gas constan t ,&:I is the he a t of so lut ion and K is the equilibrium oons t ant . S ince, as has been expla i ne d , the solubility is the measure of t he e "'1uilibriur.a constant we may replace K by M where M is the s olub ility in mo l es per 1000 grams of wa ter. The value ofAH will then be in calories per mole of solute. -128- When the comp ound under c ons ideration forms a hydra te K represents the equilibrium constant o f the rea ction ( X <a q .J) ( H2 0 (sat. s ol.)) {Xo H2 o (So l id>) Si nce the a ctivit y of t he wa t er in the s a tura t ed solution is not very d ifferent from unity in most oase s and sinoe the logarithm of K is used in the c alculati ons the error i n troduced by using the solu~ bility va lue i n stead of K is s mall . In the oaloul a- tion of the heat of solution of alloxan t h is approxi mation will be made . Over a s hort te mp erature interval 4-, li may be assumed to remai n constan t. Substituting M for K in t h e van't Hoff e q uat i on and integrating between temperatures Ti and :fl' 2 we obtain 4H 2 .303R ( T2 - Tl ) . ,- T2Tl' where M i s the s olubility at the tenp era ture Tl and 1 M2 the so lubility at T2 • S inoe these are known the value of the hea t of solution~H oula ted. oan be readily oal- -129- As an illustratien we shall calculate the heat o:f solution of adenin© . The solubility of ~denine at 25°c is . 00836 moles per 1000 grams of water or M1 =::: .00836 and T1 -:::c c98. 1 . Simila.rly ~ -= . 0252 and T2 -::. 32} . l . The value for R is taken as 1.987 . Substituting these values in the integrate4 form of the vs.n•t Hoff equation and solving for~ the heat of solution of adenine comes eut to be 8,450 cal0ries per mole . The heats of solution 0£ the ether oempounds were calculated in the same way . The results are summarized in table LV . The heats of solution are useful in calculating the value of the free energy change or the equilibrium canstant at any desired temperature if this value is known at a given temperature . This holds within the temperature range fer which the heats of solution have been calculated . For this ~urpose the van't Hoff equation is employed in the same way a.a already described . -130- TABLE LV THE HEA'rs OF SO.wUTIOJJ FOR Tfili!PERATUHE RANGE 25-50°0 At\ Compound Heat of Solution Calories per Mole Adenine 8,450 Hypoxa nthine 1 , 700 Guanine 10, 330 Xanthi ne l0.7bO Uric Acid Allan toin 8, 5,.50 Alloxan 4,180 Leuoine l.46J THE EXTENT OF DISSOCIATION OF SOLUTES I N THEIR SATURATED SOLUTI ONS AT 25°c The extent of dissociation 'ef the various compounds may be ealeulated if their dissociation consta111ts are knewn . The method ef calculation will depend on the nature cof the compound under conside r ation and upon the ap:proximatio111e that can be mad.e . The fra.ctien of t he tlissocia ted meleeules o:f a :veak acid 0r base may be calculated frem Ostwald's ailutien law o( -:.. - vK 2 +·r 1Lt + v~K~ 4 whereat._ is the degree of dissociation,v the volume in liters containing ene mol sf the electrolyte and K is the dissociation oGnstant . The caloulatien Gf the extent of dissociation of uric aeid is af this type and is as follows . From its solubility v the volume in liters required to dissolve one mel of uric aeid is 4098 . The dissoeiatien censtant ie 1 . 14x10- 6 . Substituting in the above equation ~ - 4-098x1 ,7 4x10- 6 2 t- 1u( vK + . . . . . . 081 the last term under the square root sign being negligible . The fraction of urio acid that is diss0eia ted iR its saturated solutien at 25°c ie thus . 081 .Hence the undiasooiated fraction is . 919. Us i ng the s ame expression the extent of disso ciation of guan ine, xant h i ne , allan t oin and alloxan have been oa loulated. Although alloxa n ha s t wo a ci d dissociation constan ts, t he disso ciati on is so slight tha t the second dissocia ti on has been ne glected . If both a c id and bo sio g roups are pre s ent i n t he same molecule t hen t he am ount of t he und iss o oi aI te d species may be calculated from S orensens :formula ~44 ) p 1 where pis the fra ction of t h e undissociated ampholyte. The H +- oo.noent.cation of t he satura t ed so lution must be Tl1ese have been determine d. known. The fraction of undi ssoci a ted aden ine will now be c a lcula ted. 2 . 5xl6 -7 , The I-r+ of its saturated so .;.ution is Ka is l.26xl0 -10 and Kb is 1.38 clO - 10 • Subs tit \.!t i ng ·these values in t he e quo. tion tog ether with Kw lxlo- 14 1 .996 The cal c ula tion of the undissociated fraction of hypoxanthine was done in the same way . When dealing with zwitter ions Sorensen 9 s formula is mod ified to l The fraction of un dissociated -1;. auoine was calculated by use of this formula. of KB 5 . 0xl0- 5 The value of KA is 2.5xl0- 2 and of H t- 3.16x10-?. The fraotion of leucine in the zwitter ion form is therefore 7 l + 3 . 16xl0 ... ' 2.5xlo- 2 -- + l 7 .lxl0 :t.4 5 . 0xlo- 5 x3 . 16xlo- 7 .9992 The fractions of the undissooiated s:peoies of all of the compounds are sumrna~ized in table LVI. -13.4 - TABLE LVI THE FRACTION OF THE UNDISSOCIATED SP EC I ES I N THEIH SATURATED S6LUTION AT 25°c Compound pKa pKb pH of Fraction Undiss. Sa t.Sol 9 .90 9 .86 6.6 .996 Hy_ p oxan thine 8.97 12.'/ 6.4 . 997 Guanine •••• lb .90 e O 0 .999 Xanthine 7 I., ••••• • 973 Uric Aoid 5.76 ..... ••• ••• . \H9 Allantoin 8. 67 ••••• 1 •GOO Alloxan 6.7 );8.66 . . ... ••• ••• 1 , 000 Lueoine* . l. 6 4.3 6.5 .9~9 Adenine *The dissociation consta nts of leucine are the zwitter ion constants. -13 5- ACTIVITIES 1'he activities of adenine, hypoxanthine, guanine, :x:anthine, uric aoid and allantoin were assumed to be equa l to their concentrations. Si nce these compounds are only sli ghtly so l uble in water it is assumed that they do not devia t e much from behavior of perfect solutes. The activity of alloxan in its saturated s olution was determined to be .9682 (see above). The activity coefficient of 1-leucine has been estimated by Dalton and Schmidt (4) to be 1.0'7. 'flhe activity correction may then be simp ly applied by the use of the equation, a= M( where a is the activity of the solute i n the saturated solution, M is the solubility in moles per liter and y 1s the aotivity ooeffioient . -136- THE FREE ENERGY OF FORMATION OF UNDI SSOOINrED SOLUTE I N THE STANDAHD S0LUr.l'.' I 0N A'r 25°c The free energy changes i:avo.J..-ved in going f'rom t he standard f r ee e ne rgy of forma 1iion ot1 -~ue----anhydr ous solid compound to a solution in W'1 • -e.h t ~e undissoci a ted solute is at one molal activity ·1, i.LL next be co ns ide r ed. S uppose first t·ha t the solute :,d o es not f o rm a hydrate. Then the f ree energy of t he so lute in its \ s a t ura t ed solution is e q ual to t he fri,3e energy of the soli d with whio h it is in equilibrium, X sol 1 d -"? X a q_ a sa t • ; 4 F =-0 • (l ) The free energy change in this process is zero . If the solute forms a h:1dra te, of t he g eneral formula X . 3H20, t he cha.nges t o.ki ng p laoe may be repre- sented as below, where P0 represents the vapor pressure of pure water in the standard gaseous state,P 1the vaper pressure of the hydrate-anhydrous system .P1 the vapor pressure ef pure watsr in the liquid state and Ps the vapor preseure of the saturated s olution. All of the reactions are considered. to take plaoe at 25°c. By sumnnmg up all the above reactions we obtain X(s) ~ X(sat . aq . ) ; 4 F ~ 3RTLD. P 1 1Psat (2) The free energy eh8lllg-e in the transformation ezf one mol of anhydrous x~ to a saturated solutiom. in which the hydrate X. 3H2O is in equilibrium with the aqueous phase ma.y be obtained from equatien (2) 1£ the vapor pressure of the hydrate - a:mhydr@u.s traneformatien is lmewn . t ogether with the vapor pressure of the saturated solution o As~ illustration ,the free energy change in the transfer ef one mol of a l loxan to a large amount 0£ ita saturated soluti©n will be considered . It has been shown abeve that allexa.n in its saturated solution is in equilibrium with its trihydrate aa the solid phase. The vapor pressure of the 8.l'.lhydr0us - trihydrate tr$,l}.s£ermati@n. wa s found to be 21.1 mm of Hg, Md the vapor pressure of the saturated solution was found to be 22 . 18 mm . of Hg at 25°c . The free energy change in the precess Alloxanc solicl anhydrous) ~ Allaxa.R( sat . sol . ) is given by equation (2) 4 F ~ 3RTln21 . l ::= -136 Calories. 22.78 Alloxan was the on.ly compound :found to form a hydrate. We shall next pro ceed to evalua te the free energy chang e per 11101 i n the transfer trom a sa tu.rated solution to one i n which the undissociated solute ia at one molal activity at 298 . 1°K. This is given by the equa tion, where X is the activity ooeffioient, f is the f r aotion of undissociated solu ·te in the saturated solution, and M is the molality of the sa turated solution • .A s an example the free en ergy change in the trans- fer of one mo.1 of 1-leucine from i ·ts s a turated solution to one in whioh the neu tral form is at one molal activity, will be calculated. The molality of l e uci ne in its satura ted solution was found to be .16:/ . The a c tivity coeffioient in its saturated s o l ution is given by Dalton and Schn1idt as 1.07 (4) . The fraotion o f u nd issociated leuoine was found to be . 999. Substituting these values in the e quation (3), we obtain , 4F -RTlnM tp -1364 log {. l6'7 .c . 999xl.07) 1 020 Calories 7 Having the standard free energies of formation, and having calcula ted the free energies involved in going from the solid anhydrous compounds to their solutions in which the undissociated species is at one molal activity, t he f:cee energies of formation of the undissociated compounds in their standard solutions have been calculated. in t a ble .LVII. The data are given In the second coluJru.1 are given the sta nda rd free energies of formation of ea ch comp ound 6 . at 2~8.1 K. The values of Stiehler and Huffman t2) aorreoted by Huffman (45) were used. Th e t h ird colunm gives the free energy changes a t 20a . 1°K, involved in the transfer of one mole of anhydrous s olute to a saturated solution . In the next thre e colw.uns are given the rnolali i:tj,:e,s of the saturated soluti on, t he activity a oe ff ioients used, a nd the fraction of undissocia te d (o r neutr al) molecules in the saturated s olutions, all at 298.1°K. The seventh column gives the free energy changes involve d in the tra nsfer of one mole of solute from. the satura ted solution to a solution in which the undissociated solute is at 0ue mola l activity at 298.1°K. .rhe fin :3.l oolu.mn gives the sta ndard free energy 1 of formation of the undissociated solute in a (lueous -140- solution at one molal activity at 298.1°K. This was obtained by the Slli'1lmation of columns 2, 3, and 7. -14..1.- .i..VII TAB.1.,E THE FREE ENEHGIES OF FORMA TI ON I N CALOR I ES PER MOLE OF THE UND I SSOCIATED COI/iP OUNDS I N THEI H STANDARD S OLUTIONS AT 25°0 Substanoe 4.Ft l Solid} \ s,-) ~ q . sat ., M. .sat . sol. 1ta.q.sat~ Jla ~l, undiss) ~ F =- - RT lnM / •I Adenine 71,960 ••••• .00836 l. J:iypoxan thine 18,700 ... . . .0053b 0 'l.~ pY 0 4.F f (.a = l ) undiss. 2,840 74 ,800 1. .99 6 . 997 3 , 100 21 ,800 Gua nine 11,710 ••••• .0000225 1. . 999 6 , 3<40 1 8 ,0:50 '7 Xanthine -39, 310 ... . . .000108 l. .973 5,4~0 -33, 880 Urio Acid -90,130 •• • •• .0002 44 1. .9 19 4,9 80 -85, 15 1 Allantoin -106,430,,- •• • •• ~0373 1. 1 . 000 1,950 -104 ,480 Alloxan -181, 910 - 140 .. 2 .60 1 • 00 0 -550 - 18 2 ,600 Leucine -82 , 480 ci ' , ••••• ~1 67 .999 1,020 -81 ,460 i.+ 8 .'> a.- and 1- -; :, ,, ,: t ~? : <¢ . 968 1 . 07 -142- THE Fi:tEE ENERGY OF FOlli\1A 1rION OF IONS I N THEIH BTANDARD SOLUTION AT 25°c Sometimes it is preferable to consider a reaction as taking plaoe betwe en ions rather than between neutral moleoules. It is then necessary to know the free energy of f ormation of ions in their standard solutions. Let us oonsider an aoid of the general forxuula J,CH, whioh ionizes in a queous solution according to the e quation XH , -) '--a .::1 , aq .J X . --za :::1, aq .) + The free energy of ionization is g iven by the equation 4_F ::.. -RTlnK ( 4) where K is the ionization constant of the reaction. For a general case we may wr ite the free energy of ioni zation as 4-F =: -RTlnK, K2 -- Kn where K1 , K2 --Kn represent the dissociation constants. This e quation may be applied to the free energy of ion- ization of the cation, or the anion alone, or the bivalent ion eto. It is to be noticed th at the basic dissociation constants are ordina rily c aloula.-ced in suoh a manner that t hey include the d issociation of wat er. In order -143- to be applied direotly in the a bove e quation they must therefore be expressed in terms of the dissociation of the hydrogen ion. The following e,:iuations will illustra te this relationship. Suppose tbat a substance X ionizes according to the following soheme, all the reactants and products being at one molal activity X + H o -:> x •HOH -) XH ➔ -+ OH2 Water ionizes aooording to Subtracting, we obtain X+ H'~ ➔ XH 1 The ionization of bases therefore consists of the gain of protons or hydrogen, whereas the ionization of acids involves the loss of proto ns or hydrogen ions. If the dissociation constan ts are writ ten out for each reaction, it will be seen that they are related through the ionization constan t of wa ter, i.e., wi1ere Kw is the ionization con?ta nt, of water, ¾ the ordinary dissociation const ant of the base, and 1)_ .., the 1 basic dissociation const an t in terms of H~. -144- 11i . For oonvenienoe, Kw/ Kb is used rather than On this bas i s t he free energy of ionization of a base is g iven by The free energies of formation of ions in a queous solution at one molal aotivi ty at 2~f~; 1°K ma.y be derived as follows. Consider t.he ionization reaotion of the comp ound HX to give X!la ,. l un diss .aq .J ~ t; =+a<{ 9+ ~ =l J ~e know the free e nergy of ionization from e quations (4) 1 and (5) above. We also know the free ene rgy of formation of Xlla ~lundiss . a q .) 'rhe fre e energy of formation of Hta~,, is t ak en as zero. Now, the free energy of the reaction is e q_u.a l to t he of ·the standard fre e e nergies of the SU:.'11 rea ct ants less the sta ndard free e nergy of the products, henoe, 4.F 0 f (a -::o.l, ion . aq ) o .: l und iss. aq . - RTlnK = F1\a In t able LVIII the free e ne r gies of formation of ions in 0 their st andard solutions at 298.l Kare given. In the second column are given t he free energies of formation of the undissooi8ted ions in thei r standard solutions. -145- These a.re taken f rom t able LVII. The free energies of ioni zation a r e gi ·ven i n the fourth column. In the fifth co lumn are g iv en t he free ener gies of f or 1ua tion of the anions at one molal activity in aqueous so lution at 0 298 .l K. These a re obtaine d as the s um of columns 2 and 4. Similarly in the rema i ning co~wrms t he free energies of formatio n of the cations i n the standard solut i on are given. -14 6=- TABLE THE FRE:E ENERG IES LVIII OF F0R!\dATI0N OF IOm3 IN 'i'HEI R STANDAHD SOLUTIONS AT 298.1°K Compound L\F o f Ca -=- 1, undiss ~ \ pKa ;.;RTlnKa Calories Calor ies .... 0 -'\.I! f' ( a ::c.l pKw-pKb Ri:tln:K._v/Kb an ionJ Calori es 4 F o(a : l oa. ion, i Aden i ne 74 ,800 9.90 13,500 88 , 300 4.14 Calories -5~ 550 Hypoxanthine :~ l,800 8. 9'7 12 , G4 0 34 , 0 4 ) 1 . :5 -1.770 20 . 03 0 Gua nine 18,0_5 0 •••• •••• •o•• • • 3e 10 -4 , 2 3) 13 , 8 2 ) Xanthi ne .,.33,aso 7 .10 9 I 68 .J - 24 , GOJ •••• •••• Urio .A cid -85, 1.5 0 5.76 ? ,86 0 -77 , 29 ) .... Allan to in -104 , 48() 8 . 67 11,83 0 - 92 , 6.5 J Alloxan -182,600 6.70 9 , 140 - 173, 46 0 8.66 11,810 - 101 , 6., o pK~ - 68 , 230 l.60 •••• •••• .. . . J ••• Ca l ories 69,150 •• 0, • • ••• I • ••• .•·.. .... - 2 ,180 -83 ,640 pKw-PKs 1-.Leueine -8L,460 9.70 13,230 -147- SU-.i>PLEMENTARY DATA Certain data in addition to t hose thus f ar presented are re quired befor e the fi nal oaloula tions can be made . The mos t reliable values have been selected f rom the literature . The f ree evergy of formati on of gaseous oxygen a t 0 one atmosphere a nd 25 C i s by definition t aken as zero . The f ree energy of formation of ii"" i n so lution at one mola l ao tivit y and 25°c was a lso t a ken as :z.ero. Th e free etergy of format ion of l i qui d water i n its sta ndard sta te was t aken a s - 56 , 720 calories (46) . The free e ne r gy of formation of a queous arrJID.onia i n its sta nda :rd sta te was t a ken as - 6,300 calories. (47,) The s tand~rd fr e e energy of for mation of rm f a q,. 4 is t a ken as -18,930 calor ies (J.t7o, . The standard fr ee energy of forma ti on of gaseous ca rbon d ioxide wa s taken as -94,240 ca lories {4W). From this t he standard free e nergy of formation of aqueous carboni c ac i d at 25°c was oa l culated according to Lewis and Randall (3'i)) using t heir va lues for the molal 0 solub ility of ca rbon dioxi de i n water a t 25 C at a part i al :pressure at one a t moi;,p here . '.rhe result obtained ·was -149 , 150 c alor ies. The fre e e nergy of forrnat i on of solid urea was t aken as - 4? , 200 c a l ories (4;5) ) • Fol l ow.i ng the example i n Le wi s -148-- and Randall (3~) t he st anda 1--d fre e e nergy of forma ti on of urea i n ay_ueous solu·tion was calculated . The a otiv- i t y ooefficient of . 5366 , and t he s olub i l ity of 20 . 00 free molal ( ;,~} were used . 1'he s t anda rd e11ergy of e.q_ueous urea was found to be - 48 , 440 c alori es . Jfor convenie.nce the s e a .re t a bula ted bel o w. TABLE LI X SUPP LEMfil TABY DA TA Substance 4.\F 0 r <?-=-1, aq .) 2~8.1°K Calories Oxygen (gas) D 0 -56,720 - 5 , 300 NH/ H2 CO:; - -18, 930 - 14~, 150 - 48,440 ·-149- EQUILIBRJrA OF PURINES AND THEIR DEGRADATION PRODUCTS The eourse of t he degradation of purines may be formulated as 1 C5H5N5 + H20 -:::> Adenine C-5II.40N4 -t- NH3 l Hypoxan thine t 1/2 02 C5H50N5 t- H20 ~ Guanine C5fI402M4 xanthine + Nll3 C4Ii405N2 Alloxan + CO( NHg)2 Urea Thia series of reactions may be divided into six individual reactions . Complete thermodyna mical data a.re now etvailable to compute the standard free energy change i n solution for each react ion . From t he free energy cha nges t he equilibrium c onstants and the ox idation reduction potentials can be calculated. Eaeh individua l reaction will now be cons idered. 1. Adenine,(aq. ,a:1>'1I20(liq.) --::) Hypoxanth in@{ aq ,a:•1 )-4,NH3( aq) The standard free energy change in th'.ts reac t ion i.e .,- when all -150- the products and reactants are at unit activity in aqueous solution, is t he d i fference be tween the sum of t tiet.ree energies of formation of the products a nd t he reactants. Thus for the sbove reaction AF 0 29a.1: (21,800 • 6 1 300) • (74,800 • 56,720) : -2,580 calories. From the relation AF 0 = -RTlnK , the equilibrium constant for this reaction as written comes out t o be 77.5• The deamination of adenine may be written in a slightly different manner, H~ + Adenine(aqo ,a.l) t- HgO( liq.) ~ Hypoxanthine{aq . , a.1) + NH4( aq e ,ael) A F 0 298 • 1 : (-18,930 + 21,800) ... (74,800 - 56,720) : ~15,210 calories 2. Similarl y for t he reaction Guanine ( aq. , a :l) + n 2o( liq•) ~ ..6.F 0 2980 l : ( •33 ,880 - 6,300) ... ( 18 1 0.5:> ... 56, '72Q) : -1 1 610 1 rhe XE> /1thine { aq . , a :1) -+ NH3( aq •) calories equilibrium :o:ru=<l·ent f or the reaction from right to left is 12.8. For the reaction written -151- 3. The oxidat ion of hypoxanthine may be wr i tten I:-lypoxanthine {aq., » =l ) -t ½0 2 (g) -~ Xantp.ine (aq •.• a•l) 4 F029a .1 4o :: .... 33 ,sso -21.soo : -550 680 c alories~ Xanthi ne in turn may be oxid.ized to U..t"ic a cido The reaction may be wri t·ten Xanthine ( aq •• a--i} + ½O,,{ -) :::;, Ur io a oid( aq .,~ a. -l) w g 4 F 0 ~i§,1 :: - '.3 50150 -t- 33.880 : -51,27 0 calor ies. Uria aoid may be oxidized i n t wo ways Uric a cid ( aa., a::1) -t- to 2 ( ) -l-2J3 0 g ' • ·2 (li q) - ) Allantoin( aq ., a '"'l - ) -+ HzC03 ( aq o; a -- 1 ) &F 0 298 01 : •149 ;150 • 104, 480 -( ""~x56 , 620 .. a5,l60) a · •55 9 040 calori es . Uric ac i d(aq tt;a::i) + ½o (g) f 2HzO(liq.) 2 ~) .Alloxa n(ag •• a ::;l ) + Urea {~q ., ,-a=l) 4.. F o 298 • 1 7. = -48,440 --18 2600 i- 2x56 , 720 + 85 ,150 = •32,45 0 calories. The dismutati on of xanthi ne to hypoxanthi ne and uric aci d may also be considere d. in ·this connection. It may be considered a s the sum of r eacti ons in 3 and 4 ab ove and may thus be obtained from the same data . The reac•tion is 2Xanthine ( aqu 9 a:1)) Hypoxanthine (aq. ,a•lrUric aai(l(aq .,a=l) = 21 . aoo -85,150 ~2x33 ,880 = 4 , 410 c alor ies. -152 The equilibrium constant for t his reaction or the dismutation consta nt comes out to be .000588 ,. ~'ram this it may be readil y calculated that "the amo unt of xanthine which may undergo dismutation is 2 o4% o The above dismuta tion constant refers to the undis• soo i ated produ<'.3ts and reactants. In a similar way the dismutation constant for the ions may be obtained. In practice however one deals 'Ni th the t otial t1nionized and ionized participants-. This dismutation oonstant is more useful as will be seen later. Calculations have there- fore been made of the free energies of formation of the total participants in a m.anner similar to that in table LVII • Vie thus have for the re action 2 Xantihine(total) --:1 flypoxanthine(total) + Uric a cid(total) .. 21, 800 ~ 4 ;400 calorie s~ - 85 , 200 + 2x33 , 100 The result is essentially the same as beforeo The dis- muta t ion oonat ant comes out to be .00059 5 . s. It is useful to express the results of 3 and 4 abovE:3 in terms of t he oxidation re a uction potenti also In accor- dance wi th the usage adopted in bioche mical literature, all the oxidation reduction r e ac•tion s a.re wr itten in the direction Reduotant -~ Oxidant. Suppose we c onsid.e;I" f irs t the hypoxant;hine .. xant;hine half cell reaction, wh ich may then be written as -153.:. From the standard frae energy valll.Bs we obtain - l , 040 ealories . ... •,.1 ' The value of the molal eleot1"'ode potential may then be obtained from t he relation = where 11 i s the number of electroohemioal equivalents involved in the reaction; £i' i s tlrn value of' f'arad~y wh ioh is t aken as 23 ~068 calories pe r volt per electrochemical equivalent and E0 the molal electrode potentialo ~or the neutral molecules react i ng in the direction a s written above w~ have 1 1040 E0 .0225 volts . 2x23,068 Si mi larly we may- consider the anions to take part in the reaction Yi.anthi ne - (aq. , e.;;1) + 2H1 a::1) + 2e - ; ) liyp oxanthine -- (ag_. , a =l) -I- BzO (liq .. ) = - ... 1 , 520 2:x:23,068 00329 vol ts . In t he above r ee.ctions the aoti vi ty of Ii molal. is one Experimentally however the potentials may be de - termimed in solutions at any pH . The calculat ion of Ei at pH 7 is as follows , Y.anthine ( aq ~, a =l} -t- 2fi (a=l) i 2e - ~ Hypoxa 1thine ( a g_ ,q a =l) -f ttaO (li q O ) 1 , 040 calori es. -154- lso -... -... 19,096 calories Adding t hese reactions re obtain Xanthine (aq. ,a- 1) -\- 2Ht a • 10-7) t 2 8 -> Hypoxanthin (aq . ,a=l) -t- H20 (liq.) ~F • 298 1 E0' - - 1,040 - 18 056 2x2~ . ~~e + 19,096 18,056 calories . -. 391 volts at pl 7. : Similarly for the reaction involving anions . • i' 0' - -2 2x -.447 volts at pH 7. We shall next consider t 1· e xanthine- urio acid half cell reaction. The oaloulati ons i nv olved are similar to t hose above . For t he reaction involvil'l8 the undissociated species we have Uric acid(aq. , a=l) + 2ii (a=l ) + 2e -~ Xanthine(aq.,a=l) + H20(liq.) Correspondi = - 5,450 calories. to tti.is • .118 volts . 2x~3,068 Similarly for the reaction i :"volving anions 0 C +5 450 Uric acid( aq(,,a=l) + 2H{ a=l) + 2e Xanthi ne{ aq.,a• l) + H20(11q.) 4 FO : , 77,290 - 56 , 720 - = 24 , 200 :: - 3,630 calori es . .079 volts . Now for t he correspondi g reactions taking place at pH 7 -155'- ·' 2H+Urie ac1d(aqQ~a:::l) ~ (a=io- 7) + 2e ~ Xanthine{aq .. ,a•l) + H20(lil!•) -... E' 0 -... ...132x23,068 , 646 -... ••295 volts at pH7 • also for the ionic reaction Uric ao:.i.d""(aq.,~a=l) + 2H(e.=lo ... 7fa -::> Xanthine-. (aq .. ta•l) -t- H20 (liq ) 15,466 caloriesQ =i - .335 volts at pH 7. Experimenta.lly it is more convenient to measure the potential of the hypoxanthine-uJ::-ic aoi d i nstead of the hypoxanthine .....xanthine system;, The oxidation 1:-edu.at ion potentials may be readily derived from the ab ove datao Thus we ha ve Ur ic ac i d(aq .. ,a=l) +- 4H (a=l) -f- 4 e - 0 H.ypoxan thi11e (aqcqa=l ) -6.490 calories -... .. E'0 ... . 8763 voltso - .343 volts a·t pH 7 And for the reaction involving ions £'.\ll'o E0 E'0 ... - -~,110 calories. .. ... .0228 volts o ::: ~391 volts at pH 7Q + 2HzO(liq,) -156..; DISCUSSION The reaotions for which the sta11dard free energy changes have been de.rive d above may be divided into three Adenine and guanine undergo a hydrolytie deamination to yield hypoxan'l;h_ine and xant Line .respectively. E'rom the magnitude and the sign of the free energy ehanges it may be concl11ded that these reaetions may proceed spontfiine- ouslyo On the other hand the amina ti on Qf llypoxanth+ne to adenine or of xanthine ~o guanine \JO\ild not take plaoe to any extent~ The synthesis of adenine and of guanine in vivo is probably coupled with another energy yielding reaction or it takes ~ differe.nt route to accomplish trds purpose . In t he second category are the oxidative reaetions of hypoxanthine to xanthi.ne s.nd of xanthj,ne to uric acid. These reactions have been shown to proceed either aerobically or anaerobically(49). Wh en written in such a manner as to involve molect1la.r oxygen the energetic relationships indicate that the equilibrium is in favor of complete ox.i• dation and that the oxi dative reaotio.ns are irreversible. On the other hand if the anaerobic process invol 'Iri ng hy... drolytic oxidation is considered a different picture is obtained.o This will be furt r~ er discussed in conneotion with the oxidation :.reduetion po tentials. The further degradati on of urie ac i d involves a simul taneotrn or consecutive hydrol ysis and o:xiclation. -157- In more alkaline solutions uric acid is known to be bl'"oken _down to alla.ntoin and caz:•bon dioxide~ Allantoin is tha final metabolio product; of: purine me·tabolism in many animals ~ It has also been found in plants . In an acid medium uric acid can 'be broken cl.own to alloxan and urea . JUloxan is only occasionally fo11.nd in biologioal materialEh From the s tano.ard free energy changes in the uric acid.,. allantoi.n ana. ·the uric acid-allo:xan degradation i~ aqu.aous solut:tons it may he seen t hat the equilibrium lies in favor of practically complete degra dativno These r eactions may therefore be considered irreversible unless aouple a with ot her energy yielding reactionsCI Another type of reaction whioh has briefly been consia.e.red. is the a.ismutation of xanthine to yield hypo ... :x:anthino , and urio acid ,· This has been shovm to take place (50} (51) (52). :B'rom the a.ismutation constant vie tlave calculated that the amount of xanth ina which may undergo dismu.tation is 2 . 4%,. This figure is somewhat s malle1.. than ·those reported but is significant. fhe reverse reaction 1 has also been observed experi r;ientally. The di.smutation reaction is a complioating factor in the calculation of oxid a tion reduction potentials. This reaction may also be of i mportance in the i n vivo synthesis of pur inEHh The oxia.ation reduction potentials of hypoxant.l:.dne- xanthine - uric ac i d systems has been determined experi mentally , These will be compared with the results derived from free energy data , -158~ Before mak i ng a c ompari~Jon of' our .res ults with those i n the literature it is essenti - to understand cle~rly the salie nt features involve d in each of the methods em.,. ployed by the different investiga tors and t he basis to which t heir results refer , It is furt her nec e ssary to find a common basis on whi ch the results derived b y dif• f e1"ent me thods may be compared. We shall 0011sider briell!ly the three di.fferent methods as exemplifiecl firs t, by our method based on free energy data and r epresented by the work outlined i n this thesis; seoond 9 the me thod employed by Filit ti(20;53 ; 54 } ba sed on potentiometrie meaSL1rements; arid last ,. t hat empl oye d. by Green (51) v.tJ:1ich t hough e xper imentall y the same as Filitti's cli ff e.rs i n the method and the basis of calcula tion . The derivation of electrode potentials from fre e ener :Y dat~ has alreac1.Y been illustra te d in the calcu... lations given above . The fund amental relati on be tween t he .free energy change i n an oxiclation reduction react ion and the electrod.e poten tial is gi ven by (1) AF where cl refers to the free energy c hange of the ele ctrode reaction sp ecifi ed and E r epresents the voltag e of tbe corres ponding cell. When the e leetrode equation is 1.. efer.r e d to the standard hydrogen electrode po tential the po te ntial of the cell is represent ed by Eh hen ce L\ F : -- Eh lTh" ( 2) When all the l"eaotants and procluota in t ile electrode equation are at unit activity a s for exarn.ple when the etan<l ard f:r.ee energy of the r e a ction is lrnown then ..... ( 3) where !Q 0 is the standard molal electrode potential. He1,e H-t- is at unit ac1;ivity. It is sometimes necessary- to calculate the electrod.e pot ential a.t some other H+ ac-tivi ty . There w:i.ll therefore be included in equa.tion(3) an additional term repres~nti ng the free energy c hange involved in diluting H+- from unit aativtty to the specified aotivity~ This is given by (4 ) The value of the eleotroda potent ial at a g iven pH when tl1e logarithm. of ratio of o:xidtrnt to red.nctant is zero is represented by E6• 11his is sometimes r0f0r.recl to s.s the apparent molal potential • Comb ining equa·1;1ons (3) and (4) • a gene1•a l eg_ua t.10n re l a ti• ng llio • E" o• ., and I-_i twe obt a1n -... Eo i ~.n [H-tj {5) .i! In equations (3) and (5) the ratio of oxHl.ant to red.uctant is such that the logar ii.hm of this ratio is zero . If it is required to calculate the potential of an oxidation reduction reaction corresponding to any r atio of oxidant to red.uctant at any pH a ge neral equa tion must include a term represent ing the dilution or conoen... tration of t he oxidant an d t rie reductant from unit e.ctivity -160..; to t.he desi1.. e d activity~ A F This is given by -- RT ln LRed] (6) foi j Inoludi11g this term. in equation (5) we obtain :re 0 .-.. RT ln LRedl + RTln[ n:J nJ LOX ) F (7) In ord.01.. to avoid c onfus :lon w:1e :, the poten-tial of any oxidation reduction SJrstam is given the corresponding electrotle equation should also be gl ven ina.icating clearly the species :referre-d to as well as tte aondi tions~ From the free energy dat, a l!le know direetly the standard. frea energy of :format ion of the various spe-0ies of oxidant and reduetant 411 One ean t h e1~efoi.. e l?eadily write t; he reactions eorres·ponding t o an,y of the above equations. The method of deriving t 11e fundamental constants Eo or :3b for the oxidat ion reduction reactions from po~ tentiom.e trio data will be oons ide.red nex t~ The general eq'J.at ion rela.tin,3 the observed potential to t he rat io of ox idant and reductant and to t il e Hi- aoti vi'l;y has alre a dy been given in (7) above. d ifferent~. The d a ta employed howevr;3r axe Instead of being g iven L'-¥0 o~ Q J? and calcu- lating the values of E0 ~ Eb• Eh etc-.,. the procedure is reve1"sed~ Her e from the obser ved Eh values, the :pH and the ratio of the oxidant to reduetant the E0 and tl1e ,6 values for the given system. are c al cula·t.ed .. The electrode r.eaetion is generally writeen in such a. way as ·to ref er to a singl.e type of oxidant and red.uot ant e.g •• the un<lissociaf;ed molecule the conoentratio.n -161- o:f whioh is not directly measurab le~ The oo.ncen tr a t:Lon of' the undissociated oxidant or.• recluctant according to this scheme is tJ1erofore clepende11t on t t,e di asociat ion constants and on the pli of ·the solt1ti on. Equat ions must t.he.re:fore be d eveloped to include the se variable s as 11 ell as to ex- a.re readily me-asu.rahle. The :following electr xle eq 11a.tion :f"'o.r the xru-1thineuric acid syste m is m0<lified f'1"om. F ilit;t i 's.. O:nl y one d iesoeiaj:li on c onstant is used :E'ol." ea.cl1 oonp ound!l :rhe oxidant aua. reduet-a:nt are considered to be tl e undissocia- ted s- eoies. l\riefly., t hen~ for uric acitl. -we 'ia ve tJli ~::> U"" + H ➔ from wb.ieh also UH total : and combini ng these tJ.H + U- UH Si milarly for xanthine XH - Xllt:" • KxH.,..+ H" Subs ti tut .ing these value s of till and xa for [ Cx3and £.<ted.J i n e quation (17} we obtain ( 8} -162 - This equation is in a more usef ul. f0rro. t h.an equa tion {7 sinee the oxidant and redu.otan t are expressed i n readily values may t hen be sub- messu.rab:le tern-J.S:··· Experirne11t al stituted for the various t ::-ms an-d BV- thus evalu.-at.ed f o-r the given systeni. Similar equations may be derived for any other oxidation re duction sy filtem-• iJYe shall finally eonside r t ha method used b y Greett for the evaluation of electrode potentials. He itetermined the observed potential, the pH and t he canc;entrations of t he total o . ·i dant and redt1eta nt . When t he rat io of t 11e total oxi dant to total re duc t&nt is unity the observed potential Eh t hen beoomas E; at a g iven ;pE . Th is rt 6 is not t h e sara..e as t hat derived fo r t.ne dissociated or t lle unc'U s soeia.ted s peo ies of t he ox idant; and re due taut~ Its relation t o the ele-et.r·ode a q ua t ion will be mad.a cle-ar by referring agai n to e quation (8 ),. It will be r eoalle-d that tbis eq11ation was dar i ve d for t he t1nd .isaociated s peoies of the ox idant and re ductant. ~•urther-more t he seeond and t hird terms in t he equa tion were derived b y the expansion of tha second term of equation {7). If t he 1niti o of t tJe tot a l ox idant to t he tota l re• ductant is unity equation {8) becomes - ... { 9) where E~ refers to th-e system i nvol vi ng t he t tal oxi dant ~:nd t he t ot al re duetant and E0 r epr@ sents t he mo1al eleetrode p otential f or t he und.i ss ooiated syste m•. -163- If tirn 1.. atio of the tot:al oxidant to the tota l ra- ductant is W1ity eq uation (8) become$ (9) where ·r6 refers to t he ay: tem in.vol ving the total oxi dant and the total reduotant and E0 represents the molal elec- trode pi;tential for the undi ssociated system. The symbol g~ as used b9 Green has been inter:pNted to be that defined by eqaation (9}. it possible to evaluat-e E 0 This equati©n makes for th-e unoi seooiated system ana_ thus provides a oom.t11on basis on w.l:lioh Greenfs results may be compa.."'.'ed with ours a.ml with those of Fili.tti. In a similar way equations m1-'.. \ y he derived :for t he evaluation of -eleet1rode po-trentials :ror other eyst.ems. In t h is oonneetion we s l1all reealeu.late Green 's data . Green has determined e :x-J.1eri mental.ly the E~ val ues for the hypoxanthine-:xant taine and xanti h ine-urio aoid sys:. tems at 30°G and pH 7. Re g ives the E6 valuea of - . 371 and - . 361 volts respeoti vely for these _t IYO systemt:h ·Using Gre-0n 's values for E6 and our results for t l1{z , pH and temperatu.tte disaoaiation constants end su'b- sti tuting these in appropriate equat i ons eo1"respondi :ng to equat ion {9} the :following results ~tere obtained., At 3-00C for the hypoxan.thine-xanthina system the E0 value referred to t 1:1e undissooia tea. oxi dant and reduotant <.10I:!l.8S out to be .. 057 volts ,. . Thi s corresponds to the free en.argy ohange of - 29630 calories. In a similar way E0 for the xa.nthine urio aci d system refa 1:-red to ·t he ttndissoo ia.ted. speoi es is . 090 volt.s ., the c orres.p onding standard free -164- on01.•gy change . nein.g- 4,150 oa lo r ies . t 1.\fo systems the molal I;y s ummation of tl1ese alectroa.e p-o tential E0 f or the hypo• xanthin@-urie acid system cones out to b e .• G735 vclt13w.t t h a corresponding standard. f ree energy cha nge of- 6 1 780 For ti:i.e pu.t~pone of a direct comparison of' our results with those of Green and. Fili tti we have recalaulated our v aJ.u.es for t he electrode potent i a l.s at 30 ano. 3s 0 c . The essential data are g i ve .n i n t able l, I X. . The val1J.es of ·the heat contents were ·taltan from Stiehler anc. B.nffman ( 2\} and the heats of rwl ution fr om table LV. The following example for hypoxanthi ne t'd ll illustrate t h e me th od of ealeula t ion i n vol ve·d,. 5C + 4li4 fN -+ 0 - ) By ( s ); Ey(s) ~ Hy (a1 ) Addi ng • =- •27 ., 630 4 1-1 - · 1 . 780 ~ .. 5C +- 4H 4 4 N -t 0 The 4.f.1 so obtai ned .re»res~mts the heat content of the aqne ou.s re:actant a t one molal act ivity. 11.he diffe.:t•ence between the neat conte nts of the protlua t s a.ml the react an t s g ives th e b.e a t of th e reaotion. Thi.s may be ustl d. directly i n t he van 't Hoff equation togetber wi th t he free energy trban5e of t h e react ion at a c:; iven terlll)e.i.'."'ature to cal eulat e the fr ee energy chsnge of that reac ti on at anot her temperature. From t he free e rmrgy ehanges a t 30 and 38° caloula ted i n this nmnner , t he E0 values at these two temperatures wa1"e calculated i n the same way a s be~o.rt:::i ... A comparison of our r esults wi th those of Green and Fitl tti f or the different system.a is suu".ma.rized i n table LX. I \'...O LQ I rl T.l BLE Lll !!!At Cont ent Data .1;n y$lo.t:~M 12e~ M~...t •27.630 10,760 7.780 '"'l'=0,950 •81,050 •19,850 4.11( l ~ • l!' 71-:-f· t.·) Hypoxanthine •91 , 810 8,030 AH Solution Xanthine --1-1.a ,$80 ,~.H.298 . 1 ( S ) Urio Aoid •68,i 310 Compound Water q- I © tO r-1 ·' System xa.nthina Hnoxan ·thine \'Uno.·" 1 SEh ) {Undies . ) Xa.nthine Urie Aoid Hypoxanthine Uric Aoid (Undiss.) TA:BLE LX .057 .u, .920 Jlree Energy Data E-, ( Vol 'GS) at 30° Green .090 . -0 69 Corr.- Di sm. .073 E~ (volts) at 330 .013 Unoorr . 11115 . 017 Free Ene rgy Dat;a Fillitt1 .009 .123 .066 Corr . Diam• .127 .0&8 -167- In maki nfs t he above e om.par ls on seve1:al obae.l?vationa ~:n e1""gy tla t a wo ul d :tea6. to a maximum error o~ at)out -.026 vol tg i n the ;; 0 values.- In applying the free e11 ergy tla ta to chentl.cal r e r.etions it :i.s l ikely that t he errors a.re not all i n t b..e ca ncel out ,., t3an1e cli reo.tion and vri.11 tlle)!afore te11d to In fact t he gooo. agreement of owr results with those of F' ilJ.a. t ti t h roughout , ru1d a t lea st f or the ,;;..J tl1at t he p13o'bable e rror in t t e f'ree enel"©~ data is less t ha n t h is. rre shou lcl f'ui"'ther like to point out that the potent:tometrtc methocl of obt aining electrode potentials is su b ject to s0ve1"al eouroes of error-. F':t•om the poten,-., tion1etric measuremen ts one is not certain whether t he po- tential measured a.etually referr: to t he reaction l)ostulated or whether it l"ef'er s to so me othe:t· .reaction. Furthermore in enzynm tic react ions one i s clepende:nt o.n the enzyme e.e ting a s a perfec t cat a l yst. On the ot her hand the free energy data ar e li mi tea. Dnly 'by the ac curacy with which detet'mi11ations c an be made .• There a ppears to be on e ot l1er oomplica.'ting fact,or, namely dismutat i o.n. The fre e energy data are derived for pure eonrp ounds and t t::.e refot"e the results obtai ned from tllem a.re i ndepe.nde.nt of dismutatio n . The cliract me asureraent of electrode :potenti a ls oJ: t i:e pu:r1ne systems are s t1b.jeet to t h i s com.plies. t ion for whi c h a o o.rrec tion mt1st be applied. I n fact Green's paper is al most en ;t irely devoted. to the , mathematical treat ment of adequately taki ng i nto aeoouat -168- the dismut a tion react:l.on iu t tle oaloula t ion of o:xi d. ation ~eduction pot ~ntials . TI'llitti on the other hand d-isr""e- gards .d iar.'lu,tati o:u i n her work mak i ng a oorreotion i n 'the add. endu:m t o her long ;paper.-o Just t o wh qt extent ci smut at ion t,akes pl ace under tl'.i,e experimental cond it ions of . e asuriUe, oxi dation reduct:t on p otent i als is not clear. nor ar-s ther·e st1.. i ctly quantitative data availabl..e 011 Lhe IL e.gni tuc.e of 0ons ·ta11t . '•he di smut at ion We have calcrnlat ed the d.i~:m1t:.,:b ion c o.u.st ant f rom f:r·e e energy data... i-1e l..s ve furthe r ro oa.laula tea. thi.s constant at 30° anti at 380 7 tho values being oOOO ,i '/ and We s hall next proceed to ,:levelop appt,opriate eqnations fos:- tile electrode potentials taking i nto account tl1e ru.sn1utettion . We shall t hen oalcul c~.t e from our molal el ectrode potenti als and the dis.mutation c onstants electrode p oten• t1 al s oorr es-ponding to ~reen 1 s -:-ind Filitti 's observed :po• tentials . Let us consider ·t he g eneral equation f or the liypo• Jl'.ant b.ine-uri c aoid oxid.a ti on reduction system 10 ) The r atio Hy/ UH may be varie d , t;he re:u-mining terms 'be i ng eon8hrn.t for a g iven pH . The d.:itnnut it'!;i on of hypoxan. thine and ur:lc aoid to form xanthJ.ne wi.ll also change t tlia r a tio and. thus chaf¥,s'e the value of Eh in the above equat i on. The effeat of rUs-nxttat ton will be to decrease bot b. hypo• xanthi ne and ur1o aci d by e qual amon.nt e whieh ma.:v be eva111... a ted from the a mount of xant hine forme d . The e qu ili brium -169- or fL:131 ratio o. the a'~i tla..nt to redue t ant wi ll. t hen be 3:t.ven by IIY - x m-~x where x is de·fine cl by ....,2 .... Km beinc; t he dis::n.tttation const~nt . Solving for x ~ e obtain Now beeomes h3.:(l,-,I3L1,) ...... ,li11;. +r!~2. -t--tJ?2 .-J!Jll _Ufi{ l-;:.~m). ~~ U.H{l•2Km} ... Ily i-V Hy2 -t- UE2 - 2HY UH(l•2Km) .fhis ra tio may now be substi tutied in eqo.ation (10) giving 1 a gener~l electrode equ,ation r,i.. iah takes into aocount t he effect of dismntatio n .. - This e quat ion is 2~-~ -B!, ln~ il,_l~--2_·~.,·m__~_m_1_ +'7-H~~2=·~ ~~l~ -P~ -~:::::'.L~H~L~l~~~ «Km. ~·~-)~ 4F UH(l•2Km) - liy +-1/ Hy 2 UH2 - Hy Uli {l-2Km) (11} Assuming that d ismutat i on t ak: :Ja plaoe ~H1uation (11) may be used to oalc11late t he elec t1"oda potentials of t he hypo- xanthina-urie acid s y stem. It is to be observed that when the ratio of hypo.xanthine to u.rio acid is unity no cor-reot:ton for di smutation is .reqaired and equa.t i on (10). m.'3.y be used direetly.. In the same way an equati on including among its variables the dismutat ion eons tant may be devel.o ped for t he electrode potent i al of t he xanthine-uria aoid system.. The ge11a1"al equati o:n foJ:> th i s syst~m has alraa::iy be an g iven i n ( 8 } a'bove . We 2 h s.ll co nsl cltH." t ·e ef:feot of dis- mutation on t ho :tat io XFI/ tJI-1 . F or eaeh mol,3cule of 1.ua:lo aoHi for m.a d by ·t he dism.utat io11 rieaot.ion t wo molecules of .xanth :i. na arc 11sed up. . !'he ra tio of' t he ox.i d , n."t to 1~eduot ~1:' !t will t l"m :i "be al tcrerl to m - 2a. T.ha val a.e of tffI +- a ! . rZ2.y be 0 011stde1"ed. i n terms of hy_poxanth:lne fo1;..rn.ad. We t hu s h;:ive f rom the value of ·t ll e d ism.u·t ri.tJ. o,n c onstan t a '! Hy or -"" The final eleetx-o'l.e equat ion fo1.. t.he xanthine-urio ac i d sye tem ·~o.ki ng into aocount t l1e d isru.L1tat ion is {12) -rue eleotz-ode potential of the hypoxanthl.ne-xant h::tne s ystem is not experimentally measured. di rectly du.e to t he fe.ot tti..a t the oxidation of hypoxan t h ine does not s? op at x a11tlline but proe·eeds to the urie acid. s tage . The oxidation reduo ... tion poten tial of 'ttli.s sys tem. i s therefore obtained indirectly from the t wo other syste-n1s wtlieb. ha ve a lready b e en. dise:u.ssed.• An ele-otr-ode equation mi gb. t however 'be developed for t his -171·- Wit h the ult. of' eqc1ation. (11) and (12} it is now po~sibl~ to make a more d et;ailed. c:~parison o:f our .r·e sults vJith t hose of G11eer1 and .f'il itti.- In tabla LX :F'illitt!'s observed 1-- otant ials for the hy poxan t hine•u.rie acid system are give11 toge-tl'le.r. with the ealeulated Eh values based on onr data. Si noe she uses a 1 to 1 ratio of hypoxanthine to urie aoi<l n o oorreetion for diamata t; ion is :ne.eea.sary1• It will be aeen that the agreement is ex)remel.y good. I.n table LXI we ha v-e ealoul a ted t.he E11 values f or t J:le xa.nt1:i. ine uric acid system by mes.na of equation (l2 ). The oorrest)onding obae,rved potentials repor-ted by Filitti are also g iven ,., Again Filitti trnes a l to 1 ratio of oxidant to 1.. e:ducta.u.t. Por t his ra.tio t l1e, eff:eot -of dismnt ati on was found to be entirely negligible. The disc1~e,p-anay betwe.en the observed potent i als and our e aloulate d res u.lt;g is somewhat larger than t.hat already g iven on P• +66~ Thia is partly due to the taet that dissociation eonstants nsed in our caJ.eulat ions wer,e different from thos-e used. by Filitti. Table LXI.I s hows (Treen• s ex-pe.r t mental value.a fo.r the hypoxanthine,• urio ac-id oompared in the saw:e way. Green uae.s sever-al different ratio s of hypox.an'thine to t:tri(l aeid. We ha ve calculated :fil-h values 001--reapo.nding to Green ts ush:ig b-oth e quation (10} and. (ll). It r.ril.l be observed t .ha. t t he potential s oalcul.at ed by mean s or equation (10) a.,.~re-e mox-e closely with Gre-eni,e da ta, us.e of equation (11}., t han those ob t aine d by the It appe a rs from this eomparison that -112--· TABLE LXI POT E IT I AL;; DERI VED 1i'Rffi.1 FIIEI: ElW}iiRGY DATA .,.,. cnloulated pII Eh observed 1 7,. 03 ... ,. 387 volt.a -~isa7 volts 7.08 -.390 -.391 .,Q8 -.378 - .• 391 7-.21 - .40 -.404 7 .,29 ......,4-10 .....407 '7 .30 ... .,410 - . 408 7.31 ... A .0 9 .... .,409 7 .• -31 -.410 -...,409 7 ..,1 ·:; .....418 -.,.4l'7 7 .• 47 - .• 420 --421 r1 ., 6 1l - ....437 - .• 436 7.61 -..J136 -.436 7.68 - .• 435 -.4 36 y Jllh -173- TABLE !JUI COM:l?ARIS01f OT!' FILITT I' S OBSERVED E.LECTRODJS J.?OTEiiTI.ALS FOR THE XAIJTIIIID3-tnUC AC I D SYSTEM AT 38°c WI TH CALCULATED pH observed Eu volts Eh cialculated volts 7 . 4,7 '""#409 - . 383 7 •.50 - .410 ... 335 '7 •.52 .... 410 •0 386 9 . 60 - .. 424 ... . 3,2 7 . 60 ..... 420 • 0392 7-.60 "" 0415 • , 392 7.62 ... .. 420 - . 393 7.64 .... 420 .... 395 7.64 • ~419 .... 395 '1 o-65 .... 420 -• 395 71!'65 ... ~424 .... 395 7 06•"' ... . 420 -, 396 7 .. 68 ... . 420 •0 397 7. 70 -.420 - • :399 8 . 35 .... . 464 ... .. 441 8, 37 ... o,463 .... 442 -174- little or no d ismutation t a kes pi ace under the conditions of Green ' s e1rperimentso Fi n ally in t8.1 Jle LXI II Green's l"'esr1lt s fo.t• the x.anthine u.r i c ac .id a re c om:pareo. with cal• cula tea. resul ·ts bau,e a. on .f1,ee energy d.a tao By the use o f eq11.1;1tion (12} ss well as equation (8) it was found that 4.t1cA. t .t e effect of 6.ismuta t ion on the xanth i ri e w:•ic- ;p otent ials i s negligibleo 'rhe agreeme nt b etween t h e measured 1.u: J.d ·t11e ealcula te6. potentials is .not satisfactory ~md. is probably outside the error in U1e free ene1•gy data. The g ood agre e• meu t of our res ults for •he hyp oxau1;hine-urio acid syst em \vi th ·those of G·.reen and. ]Pili tti mak es us inclined to believe that the free e11ergy data are reliable also fol_., ·the xan:!;hine- u.rio acid system . It must also be rem:f1ro.bered t h.a t ·the oxypurines have a limited solubility in v1a·teJ:1• and tihe mani- pulation of relatively h:l~h conoe.ntrat ion s of t h ese sub ... stances as in Green 1 s an :1 Jj':tli tti 's experimen ts intro• due es an uncertain.t y i n tlle r e s ult; s • Ii' u.r t liermore the a d- di ti on of a relatively l arge amount of the enzyme :preps.ra- t ion introduces other sutstanc~es whioh may affect tl1e el.e ctroa.e potentials.- On the whole the usefulness of thermo- dynamic data in interpreting the results obtained fr om equilibrium measurements ancl i n ;provia.ing an independent method by which the results may ·be chaoked i:1.as been il- lust1~atea.• ]'rom the va lues of t;he oxidation re cl.uotion potential s of oxypurines it nay be c ondluded th at they are mutually o:xidiza.'ble and reducible .. However the medium in which these .reactions may take plac e must be strongly re due ing. -175- T.AEL3 LXI I :t pH Hy/UH 7. 65 9/ l 8..:07 ff 8 oL.i: 3 3% obs., V•)l·l;s Eh oalo• volts E h ca.le ,., eor r. volts 1 :1 5•~k ... ., 433 .. .,4 94 !lh-8'.l: ... . 464 ... . 525 n --~4 9? .. . 490 ... 151 6• 93 4/1 .... ~382 ... "375 .......i29 7 .32 rr ... 419 •a :i:05 ... ,..{1,59 6 d'93 1 /1 .... 366 ... ,.,366. ,, 7. 32 n - ~tl-00 -. 39 ' 7.,55 n -,4 25 .... 41 9 8 ,0?' !J .... 456 ~ ..L .l:50 •• 8 . 43 !! .... 4 59 ... .. ,J.,76 fl 0 9 111 10 If .... , 497 -~ 523 •• 7 ,. 32 1/4 ...,380 - ~38 7 .... 333 7 . 65 n .... 406 -.41 0 - ~356 8 10 07 n •~432 "" *441 .,.e 38'/ a.43 !f .... 446 .... , 46 7 -«413 - .a . ... 11: ♦ TABLE LXIV COMPARI SON OF GREEN ' S OB3ERV1.J:D ELECTRODE I?OTENT I ALS :i?OR rnHE X:A.l\JTHI NE• URI O AC I D SYSTEM AT 30°0 W'I '.i.1H Oli.miCULATED l'Orl1ENT. I ALS DERI VE".D ]~ROW Flt.EE ENERGY DATA pH Hy/UH E11 obs, volts Eh calc , volts 6.93 4/1 - , 3G2 - .338 7,32 ff -♦ 39 0 - ~365 8,07 f§ .... . 452 •1t4l 6 ....465 -,439 8.43 n 6.93 1/ 1 .... 349 ""q:330 7 .,32 u ... . 380 • 0359 7~65 fl ... , 415 • , 380 a.07 n ... 440 -,407 8,43 " -.454 -.430 9.10 n .. ,.475 -~470 6093 l/4 .. , 335 -.360 '7.32 f-1 • 1> 370 -, 350 7,85 t'I .. . 400 - ,371 a_._43 rt ... . 43'1 ... ,. 420 Under ord i nary conditions o:xi.dat ion reac~ions would r>red omina te °' In conel nd:i. ng this s tud.y of equilibrium. relat i ons of pnri nes and their deg radati on produets we should like to :point out that th® standard free ene:rgy ehanges in ~queous solutions do not represent the actual conditions i n vivo, al thoo.gh certain br~a d c rnolusions may be drawn .. It is hope d. tha t the informa t ion prese nted in this thesis will find further a:pplieati on to the biochemistry of' purines - -178- SU1VIMARY The following expe riment al data for ac:lenine • hypoxanthine , guanine, xanthine, uric a ci dt allant oin 9 alloxan, and d-1- Leuoine have been presen ted: Solu• bilities at 25°C and 5o 0 c; di ssociation cons tants; vapor pressure of alloxan- a.l loxa.n trihyd.rate systemi evi dence for non existence of hydrates of aclenine and xanthi.ne; activity coefficients of alloxan, ...... all re ferred to 250c . The expe:ri men tal method.s used for the d.e ·t ermi na.• t ion of the abot_ve physical chemic al oons tan ts ha "le been clescrl bea.O' Utilizing the c:x:peri ment al clat a the stand.a.rd free energi es of .foriaation of undi ssoc ia tecl molecules II and ions in aqueot1s solution at 25° 0 have been derived; and tlle methods of oalcula ting th ese have ·been i ndi c a ted . The ane1..gy relations hips of the diff'e.rent processes involved. in t he a.egradati on o.f purines ha ve been dis ... cussed. -179-- Rfil'ERENCES 2,, Stiehler and Huff'rnan 11 J .A11C,SH §.7.~ l74li (1935) •· 3 -. Muller, Soi -u l!i.• 50f (1935)"' - 4 a Dal ton and Schmidt , J II B" C •, 103 , 549 , {1933} • • •7 ij Hossin!., Chem. Re v., li, 23 3 (.1936} •. 6~ Hunter, Bioohem. J~ ~ 16 9 640 , (1922) . 8 .• Kirk and. Schmidt , J . B. c., Bl, 237 , (1929)• 9. Schmidt a_r1d ·Hoagland , Univo of Calif . Pub. J?hysiol o, 5 . 23, (1919) ♦ - 0· lOo H:ttoheook, J_, Geni> l?hysiol •., i,, 747,. (1924). 11. International Cri tio:al Tables , Vol• III. M0Graw• Hill Book Co o, Inc .• New York anu. London., (1929) • 12~, Hunter and Hlynka, Bi ocb.em. J •~ ~, 486,. (1937) • 14. Koesselp z. Physiole Chem. , 1,Q., 254 , (1886). 1:3. F'ox tt Soi. , 84,163, (1936). us. Tafel and Ach; Ber ., 34 , 1178, (1901). 16; Bruhn.a , Physiol . Chem., li• 566, (1890), ~ z l' 17. Scherer , Ann ~, 2.Aa 33lt (18 50) • 18 o Stutzer, z. Anal. Chem • 19. Fisoher, Ber., - 30 9 · 9 .:21,, 502• (1892). 2226, (1897} • 21., Wood, J .c •s. • ~, 856. (1903) • 2la. Wood , J .c.s ., .§2., 1842 , (1906} . 22 . Strecker , Arm., 108, 143, (1858),. 22ao Strecker , Ann., ll8 t 168, (1861)0 -180- 23, Sundwielt , Z+ J.?hysio.l -• Chem , i 78 11 1847, (1912) •· 24io Almen , Jahrsberiehte, ~• 534; (1862). 25. Steadeler, Anno:1 ll!,a 35, (1859) o 2611 Horbaezewski, z~ Physiol. Chem:o, .&,~; 226 , (1897) it 27., Benedict a.~ d l!"rfu"lk.e., Jo.Bs Co, la,, 837 ~ (1922 } . 2s. His &lW- Pauli z .. Physiol. Chem •• 1, 1. (1900). 29_., Gudzent _ 1 Z 1> Physiol<A Chem.;1 §Q_, 25 , 38 , (1901) ,. 30. Young and Musgrave,, Biochem. J ., .&i• 941~ (1932) $ 3l •- lWa:rt ignon, Ami. De6himc; e 1J Phys It,. 28 • 289 0 ( 1893) ., 32.• S.chu,lze and Barbieri ,; J • I'ra.kt 11" 25 ,1 1480 33. Gortner, J ll A eCeSe;, - 33, 85e (1911)11 35. Bog ert~ J.A.c.s., y , 809, (1910), 36 "' V1heeler, J ,d:1.aC 11 So~ 3,~Q 809• (1910)• 37• Trubsbach~ Z Q Phys . Chem.~ l.po 711 :p - (1895) . - 38. Sca'Gchard, Ea:r..l:lf:r, ana. Vi'oOd p J o.A oC . Sq t 60 ~ 3061 , , , (l938to 39.- Lewis and Randall, The1.. modynamics . HcGraw•Hill Book 00 11, New Yorj: ancl Lond.oa , (19 23), 40 • Sano p Bioo hem. z.. ., 16~, 14 1 (l926) o 41., Barris, r~~o_c • Roy. Soc . E~ ,9.,[i o 4:40• (1924), 42• Win.Ir..leblech 1 z. Pllysiol, _lit 546, (l~Ol) • 43-. Wood,· J .• c.,sl) IP !9.ft11 1988 , '.1914) • 44 . SQrensen_, Ergebn . J?hys;i ol, !fi,. 393 , (1912),. 450 Huffma_ n , anpublis ho d da t~o 46~ Gia.uqu,e an d Ashley, I-'hy.s"' Re v . I) ';li,§.• lr, (1933) • 47 o Borsook and Huffman ~ Ch!'/ XV Cl1emistry of' .Amino Acids and. Proteins, Charles C. Th omas . Balt i mor e-, (l938) fl -181-· 48 o Ros s ini a.no. J 0-ssup, J . Re:3ea_rch , Wo3 os. i 1:1,, 4_9 1, (1938 ) • Stewart and I-lopld :1s, ?.roe . Eio2r ,, Soc. B.,~ 94, .... ~ - 109, (1922) ., - 50 0 Bach and hli c t lin , Ber~, 60 , 82 1 (1924) ., - 51. Gre~n i Bi oohem. J,,, 28, 1550 , (1934)., - 52. Re ino.eil a nd Schuler, Z,, J?hys i ol a Chem~ t 247, 172e (1937) • -- 53. Fili tti, 0 • ~l .. .h.cad:. Sci o 1 97 0 1212, (1933)e