CXV. THE CREATINE-CREATININE EQUILI-
`BRIUM. THE APPARENT DISSOCIATION
`CONSTANTS OF CREATINE AND
`CREATININE.
`
`By ROBERT KEITH CANNAN AND AGNES SHORE.
`From the Department of Physiology and- Biochemistry, University College,
`London.
`(Received June 30th, 1928.)
`
`THE facile conversion of creatine into creatinine under the influence of strong
`acids and the partial reversal of the reaction in neutral and in alkaline solutions
`have long been familiar. Yet only recently has there become available any
`quantitative data upon the equilibrium conditions. Hahn and Barkan [1920]
`were the first to report any systematic kinetic studies. They determined the
`equilibrium constant in solutions of sodium hydroxide of varying concen-
`tration, observed that the component velocities increased with increasing [OH-]
`and showed that the order of the reaction creatinine->creatine, under these
`conditions, was that of a reversible monomolecular system. In a molar solution
`of hydrochloric acid, on the other hand, the reverse reaction went to comple-
`tion and followed the course of a simple monomolecular change. Hahn and
`Meyer [1923] later reported a few observations which indicated that the velocity
`of this reaction in buffered solutions increased rapidly from PH 6 to 4. A more
`elaborate study of this system has been made by Edgar and his associates.
`Edgar and Wakefield [1923] determined the monomolecular velocity constants
`(k2) of the dehydration of creatine in hydrochloric acid solutions of varying
`concentration and at various temperatures. They succeeded in relating k2 to
`the temperature by means of the Arrhenius equation and, further, concluded
`that k2 was, probably, proportional to the hydrogen ion activity. Finally,
`Edgar and Shiver [1925] have made an extensive series of determinations of
`the equilibrium constant (K) at 50° in buffered solutions of PH values 1 to 6.
`Hahn and Barkan had suggested that their observations could be interpreted
`upon the assumption that the molecular species whose concentrations deter-
`mined the equilibrium were the undissociated molecules of creatine and of
`Confirming this, Edgar and Shiver obtained fairly satisfactory
`creatinine.
`agreement between the observed values of K and those calculated from the
`dissociation constants of the two bases and the value of K in unbuffered solu-
`tion (i.e. where the two reactants were not significantly dissociated).
`
`001
`
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`
`

`
`CREATINE-CREATININE EQUILIBRIUM9
`921
`
`Consideration of the results summarised above will indicate that, although
`the hypothesis of Hahn and Barkan has been useful in co-ordinating the
`equilibrium data, it fails to comprehend the relations between [H+] and the
`velocities of the reactions. If the equilibrium be determined by the ratio of
`the concentrations of the undissociated molecules of creatine and creatinine,
`the velocities of the two contributing reactions should be governed by the
`same factors. Thus, if kb be the dissociation constant of either reactant and k
`the velocity constant for its decomposition into the other, then k should vary
`[OH1]
`That is to say, the velocity should be inversely proportional
`with
`kb +[OH-]
`to [H+] on the acid side of the buffer range of kb and should be, independent
`of [H+] on the alkaline side. The available experimental evidence, however,
`indicates that the velocities are proportional to [H+] in solutions of strong
`acids and are inversely proportional in strongly alkaline solution. The reactions
`under discussion are of such direct biological interest that it was decided to
`undertake a series of kinetic studies in buffered solutions between PH 1 and
`10 as an attempt to elucidate these discrepancies.
`It was necessary, in the first place, that there should be available de-
`pendable values for the dissociation constants of creatine and creatinine. Since
`the values in the literature differ rather seriously, a redetermination of these
`constants was undertaken.
`
`DIssOCIATION CONSTANTS OF CREATINE AND CREATININE.
`The method employed was that of electrometric titration of dilute solutions
`of the two bases with standard hydrochloric acid in the presence of the
`hydrogen electrode. The routine technique of this laboratory has already been
`described [Cannan and Knight, 1927]. The reference electrode was a saturated
`calomel cell which was standardised against 0 05 M acid potassium phthalate
`[Clark, 1922]. Two palladinised gold-plated platinum electrodes were em-
`ployed as duplicate hydrogen electrodes. No difficulty was encountered in
`attaining stable potentials in any of the solutions titrated and the two elec-
`trodes agreed within 0 3 mv. at all significant points on the titration curves.
`The creatine was prepared from a good commercial sample by repeated
`recrystallisation from water. After drying to constant weight over calcium
`chloride, a typical preparation gave
`2819 %
`Water 12-18 %
`Nitrogen (Kjeldahl)
`...
`Theory for C4H9O2N3. H20
`28-19
`12-08
`A saturated solution gave no reaction for creatinine upon applying Weyl's
`test.
`Creatinine was prepared from the creatine by treating the latter with
`hydrochloric acid gas and subsequent liberation of the base by aqueous
`ammonia. The product was recrystallised from acetone [Edgar and Hine-
`gardner, 1923]. Nitrogen and water determinations were quantitative for
`anhydrous creatinine. Folin's colorimetric method for the determination of
`Bioch. xxn
`59
`
`002
`
`Harvest Trading Group - Ex. 1113
`
`

`
`922
`
`R. K. CANNAN AND A. SHORE
`
`creatinine (using creatinine picrate as standard) gave results in agreement
`with the nitrogen values, but this method is, admittedly, not sufficiently
`accurate to detect traces of impurity in creatinine.
`It will be convenient, throughout the paper, to conduct the discussion in
`terms of hydrogen ions rather than of hydroxyl ions and, consequently, all
`constants will be treated as though they were acid constants. That is to say,
`the kation of a base will be regarded as an acid which dissociates a hydrogen
`ion [Bronsted, 1923]. The constants so derived (k') are related to the familiar
`kb values by the equation Pk' = Pk,-Pkb
`
`Table I. Uncorrected apparent dissociation constant of creatinine.
`Molar
`conc.
`041
`-
`0.1
`0*02
`0.1
`
`Authors
`Wood [1903]
`McNally [1926]
`Cannan and Shore
`
`McNally [1926]
`Eadie and Hunter [1926]
`Hahn and Barkan [1920]
`Cannan and Shore
`
`0.1
`0 04
`0-02
`
`Temp.
`40*2
`40-0
`300
`30 0
`25-0
`25-0
`20-0
`17-0
`15-0
`
`Pk'
`2-97
`4-42
`4-77
`4-72
`4-78
`4-71
`4-87
`4-44
`4*91
`
`kv
`3-57 x 10-11
`1 01 x 10-9
`0-98 x 10-9
`0.15 x 10-9
`0-76 x 10-9
`0 70 x 10-9
`0-64 x 10-i
`0419 x 10-9
`047 x 10-9
`
`Wood, and Hahn and Barkan calculated kb from the degree of hydrolysis of solutions of the
`hydrochloride; Eadie and Hunter employed the electrometric titration; McNally's results are the
`mean of results from the conductance, hydrogen ion concentration and distribution of the hydro-
`chloride.
`
`In Table I are assembled several determinations Of Pk' for creatinine to-
`gether with values calculated from the kb values recorded in the literature.
`The important effect of temperature upon the constant is evident and renders
`difficult the comparison of the results of different observers. But it would
`seem that, apart from the two earliest determinations which were made with
`methods open to considerable experimental errors, the various values are in
`substantial agreement. If is unnecessary, therefore, to report our experimental
`data in any greater detail. For purposes of the analysis of the kinetic studies
`which follow, the value for the dissociation constant of creatinine at 300 will
`be taken to be k' = 1-90 x 10-, i.e. Pk, = 4-72.
`The case of creatine is less satisfactory. The various determinations are
`summarised in Table II. In Table III is given the analysis of a typical titration
`curve to indicate the degree of concordance of the data. The calculations have
`been made with the aid of the Henderson-Hasselbalch equation. The values
`of [H+], used in calculating the "corrected equivalents of acid," are obtained
`from the observed PH after correction for the activity of the hydrogen ion by
`the equation log 'TH= 0 20 V-ivz [Simms, 1926]. TH is the activity coefficient
`ratio for the hydrogen ion, Xivz is the sum of all the ion concentrations each
`multiplied by the z power of its valency. The value of z was assumed to be
`unity. The constants have not been corrected for activity.
`
`003
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`Harvest Trading Group - Ex. 1113
`
`

`
`CREATINE-CREATININE EQUILIBRIUM
`
`923
`
`Table II. Uncorrected apparent dissociation constants of creatine.
`Molar
`conc.
`0.1
`0.1
`0-02
`0.1
`0.05
`0 04
`0.05
`
`k',
`2-1 x 10-3
`2-4 x 1O-3
`2*4 x 1O-3
`2-2 x 1O-3
`09 x 1O-3
`1-4 x 1O-3
`2-45 x 10-3
`
`Pk'l
`2*68
`2-62
`2*62
`2-66
`305
`2 85
`2*61
`
`Authors
`Wood [1903]
`Cannan and Shore
`,,
`9,,
`Eadie and Hunter [1926]
`Hahn and Barkan [1920]
`Cannan and Shore
`
`Temp.
`402
`30 0
`30 0
`25X0
`20-0
`17-0
`17-0
`
`a
`- a
`
`log
`l
`
`Table III. Titration of 50 cc. 002 M creatine with 041 M hydrochloric acid.
`Corrected
`equiv. acid
`[HCI] -[H+]
`[H+]
`corrected
`[creatine]
`Titre
`PH
`Pk'
`0*00
`5-77
`000
`-
`-
`2-67
`-1-74
`0.0181
`39*0 x 10-6
`4-41
`0*20
`3*95
`2-62
`1*33
`0 0452
`11*2 x 10-6
`0-52
`23*4 x lO-6
`1-02
`3-64
`2-63
`1.01
`0-0901
`51*3 x 10-6
`2*62
`0 68
`0*1744
`3*30
`2-01
`2*63
`85-1 x 10-6
`3-02
`02568
`3 08
`0-47
`2*62
`123-0 x 10-6
`2-92
`0-30
`0-3337
`4-00
`166-0 x 10-6
`2*63
`0*16
`0 4064
`2-79
`4.98
`224*0 x 1o-6
`2-67
`-004
`2-63
`0-4749
`6 00
`3470 x 10-6
`2-48
`0-6010
`8 02
`2-64
`+0-18
`2*70
`10.00
`501*0 x 10-6
`2-34
`0 36
`0 7000
`2-76
`7940 x 10-6
`0-61
`0-8016
`2-15
`13*03
`It is probable that the last two calculations suffer by reason of the uncertainty of the correction
`for hydrogen ion activity.
`It will be seen from Table II that differences exist between the determina-
`tions of different observers which cannot be attributed to differences of
`temperature or of concentration. In particular, it is difficult to explain the
`conflicting results of Eadie and Hunter and of ourselves since the same method
`was employed and was prosecuted with the same degree of precision. No
`plausible source of error in the titrimetric method peculiar to creatine suggests
`itself. The possibility of a significant amount of conversion of creatine into
`creatinine during the course of a titration seems to be excluded by the velocity
`measurements recorded in the second part of this paper. Provisionally we
`will take the value for k' at 300 as 2-40 x 10-3, i.e. Pk' = 2.62.
`A question of some interest to the chemical behaviour of creatine arises
`from a consideration of its electrolytic dissociation. The conventional formula
`for creatine contains both a carboxyl and an amino-group. Creatine might
`be expected to behave, therefore, as an ampholyte. Only basic properties
`are, however, evident in its chemical behaviour and only one dissociation
`constant is detected by titration. This is, therefore, described as a basic
`constant. Hahn and Fasold [1925] have, however, found that the solubility
`of creatine in solutions of sodium hydroxide is greater than in water and they
`conclude that some dissociation of creatine as an acid occurs in solutions of
`great hydroxyl concentration. From their observations they calculate a value
`of 14-28 for pkai Now the allocation of the first dissociation constant of
`59-2
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`004
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`

`
`924
`
`R. K. CANNAN AND A. SHORE
`
`creatine to the amino-group and the assignment of only negligible acid pro-
`perties to the carboxyl is difficult to justify upon the grounds of organic
`chemical experience. Yet it is usual to describe creatine as a base. A more
`plausible interpretation of the acid-base behaviour of this substance would
`seem to follow the application to it of Bjerrum's [1923] treatment of the amino-
`acids. The first constant (k1'= kIW/lb) then becomes the acidic constant and
`the second-in this case, inaccessible-constant (k2') is the association con-
`stant of the basic group. With this assignment of constants creatine becomes
`an acid comparable in strength with other carboxylic acids. At the same time
`a new difficulty is created for it is required that the basic dissociation shall be
`as great as that of the alkali hydroxides. It would be difficult to concede this
`to a simple amino-group and it is of interest, therefore, that creatine does not
`behave as a primary amine either towards nitrous acid or towards formalde-
`hyde. In this connection the strong basic properties and anomalous behaviour
`of guanidine itself will be recalled. It is significant that several of the structural
`formulae which have been proposed to explain the anomalous reaction with
`nitrous acid contain a nitrogenous group which might be expected to dissociate
`strongly as a base [Hunter, 1928, p. 99].
`The above considerations in no way prejudice the application of the
`dissociation constants to the co-ordination of kinetic data. It is a matter of
`no immediate moment whether the velocity of dehydration of creatine is
`determined by the concentration of undissociated creatine or of " zwitterion"
`-the mathematical relation to kl' remains unmodified.
`CREATINE-CREATININE EQUILIBRIUM.
`Solutions of creatine (0-0106 M) and of creatinine (0.00354 M) were pre-
`pared in a. series of the 005 M buffers recommended by Clark [1922]. The
`mixtures were covered with 10 cc. of toluene and stored in stoppered bottles
`in an air-bath maintained at 300 ± 10. At intervals appropriate to each
`experiment, a sample was removed and the concentration of creatinine present
`was determined by the method of Folin. The standard solutions for this
`method were prepared from a purified specimen of creatinine picrate. The PH
`values of the various reaction mixtures were determined at the beginning,
`and again at the conclusion, of each experiment by means of the hydrogen
`electrode. At the end of each experiment determination was also made of the
`total creatine + creatinine. In agreement with other investigators it was found
`that some conversion occurred of these two substances into products which
`reacted neither as creatine nor as creatinine. The extent of the loss during
`the period of experiment varied from 0-5 to 5 % according to the PH of the
`solution. In view of the temperature at which the solutions were maintained
`and of the precarious antiseptic properties of toluene over long periods, the
`occurrence of bacterial decomposition may be suspected. This source of error
`cannot be absolutely excluded but the results are so concordant amongst
`themselves and fit in so well with the equilibrium data of Edgar and Shiver
`
`005
`
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`

`
`925
`
`10
`100
`
`3
`
`._
`
`4
`
`5
`
`6
`
`7
`
`8
`
`9
`
`10
`
`CREATINE-CREATININE EQUILIBRIUM
`(obtained at temperatures from 25 to 1000) and the relation of the losses to PH
`is such that we are persuaded that the observed destruction of reactants was
`not due to bacteria but to irreversible chemical decomposition to an extent
`similar to that recorded by the earlier observers. In any case these irreversible
`changes are not sufficient to explain the gross relations between the velocities
`and pH which will be established.
`Fig. 1 presents a summary of one series of observations upon the two
`reactions. It presents several unexpected relations of the velocities to PHu
`notably the well-defined pu optima.
`PH
`2
`
`O
`
`90
`
`80
`
`1903
`
`1.362
`
`70X
`842/
`11acanrmiistf30
`.~60
`
`1718
`
`00 385
`
`%
`
`40
`0
`
`30
`
`20Z-
`
`1~~~~70
`
`12
`
`10
`
`...'
`
`25
`
`6
`
`7
`
`8
`
`9
`
`10
`
`1
`
`3
`
`4
`
`2
`5
`Fig. 1.
`pH
`*Percentage creatinine formed in solutions of creatine at times (hours) indicated.
`o Percentage creatinine remaining in solutions of creatine.
`The data upon the change creatine-+creatinine are more extensive than
`those for the reverse reaction and the analysis of the former will suffice to
`bring out all the important relations. The velocity constants for the latter
`reaction, in the PH range where it is significant, fully confirm these relations.
`In solutions acid to about PE 3 the conversion of creatine into creatinine
`is seen to be substantially irreversible. In such solutions the reaction would
`be expected, therefore, to proceed as one of the first order. The first half of
`Table IV indicates the measure of constancy of the monomolecular velocity
`constants (I2) derived from a typical experiment and the slight effect of
`allowing for the reverse reaction. The results are as satisfactory as could be
`expected in view of the limits of error of the colorimetric method, the slight
`
`006
`
`Harvest Trading Group - Ex. 1113
`
`

`
`926
`
`R K. CANNAN AND A. SHORE
`
`changes in PH accompanying the reaction and the simultaneous irreversible
`destruction of creatine and creatinine. In solutions alkaline to PH 3 the
`system is significantly reversible. The appropriate velocity equation may be
`put in the form
`
`l + k2=IKl )
`...... (1),
`where k1 is the monomolecular velocity constant for the hydration of creatinine,
`we2is the monomolecular velocity constant for the dehydration of creatine,
`K= k2/kl, while a and x have their usual significance.
`If K be known, this equation may be solved for k, and k2. Edgar and
`Shiver give the following relation for K (when the reactants are not measurably
`dissociated, i.e. in unbuffered solution)
`-1084
`log K= T-84 + 3-3652.
`When this is solved for a temperature of 300 the value of K is given as
`0-6125. According to the same authors K is related to [H+] by an equation
`which (yvhen k' values are substituted for klb values) takes the form
`...........(2),
`
`K=06125k"[k'±[H+]]
`K = o.6125 kk[k" +[H+]]
`where k' is the dissociation constant of creatinine and = 1(90 x 10-5, and k
`is the dissociation constant of creatine = 2-40 x 10-3. From equation (2) we
`have calculated the values of K at the various PH values of our reaction
`mixtures. When these are inserted in equation (1) together with the corre-
`sponding velocity data, the term kI + k2 is found to be reasonably constant
`within a single velocity experiment. One such result is summarised in the
`second part of Table IV, in which values of kI + k2 are contrasted with the
`values of k2 calculated as a monomolecular velocity constant. Finally, k, and k2
`have been calculated for each experimental PH and the results are assembled
`in Table V. This Table is restricted to the same series of experiments as are
`shown in Fig. 1, while in Fig. 2 the values of Ic and k2 have been derived from
`two series of observations on the rate of dehydration of creatine and one series
`on the rate of hydration of creatinine. The curves have been further extended
`into the extremes of acidity and alkalinity by; the rough calculation of the
`values of kI and k2 from the observations of Edgar and Wakefield and of Hahn
`and Barkan respectively. These involve an uncertain temperature correction
`ntid can only be regarded as approximate.
`Restricting further discussion to the range of PH covered by our own
`observations the chief point of interest is that although equation (2) satisfies
`the equilibrium data it is not adequate to define the velocity relations. That
`is to say, the individual velocity constants display a relation to [H+] which
`is not apparent in the equilibrium constant. The particular relation is the
`retardation of both velocities alkaline to PH 3. Since this is not reflected in
`a change in the equilibrium constants the 'factors responsible must have the
`same influence upon the two reactions. Indeed, it may be surmised that the
`same factor is responsible for the changes in both Ic and k2 within this range.
`
`007
`
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`

`
`CREATINE-CREATININE EQUILIBRIUM
`
`927
`
`Table IV. Rate of dehydration of creatine (0-0106 M) 30°.
`
`a
`100
`
`100
`
`25
`75
`125
`170
`385
`865
`1346
`2017
`
`25
`50
`75
`125
`170
`385
`695
`1130
`1896
`
`x
`
`1-97
`5-58
`9-12
`12*25
`24-33
`46-18
`59*25
`76*92
`
`7*29
`15*39
`21*77
`31-93
`38-17
`64-62
`79-98
`83*33
`85'00
`
`k
`
`[creatinine]
`[creatine]
`tin a
`80-0 x 10-5
`76-4 x 10-5
`76-4 x 10-5
`76*8 x 10-5
`72-5 x 10-i
`71-5 x 10-i
`66-7 x 10-5
`72-7 x 10-5
`pH3-77: K =5-673.
`304*0 x 10-'
`334*0 x 10-5
`327-0 x 10-5
`308-0 x 10-5
`283*0 x 10-5
`269-0 x 1o-5
`232-0 x 10-5
`159*0 x 10-5
`99.0 x 10-5
`Table V.
`
`Ka( l)
`kL+k2= In
`80-0 x 10-5
`76*4 x 10-5
`76-4 x 10-5
`76*8 x 10-5
`73'1 x 10-5
`72-9 x 10-6
`68-3 x 10-5
`75-2 x 10-5
`
`350-0 x 10-5
`398-0 x 10-5
`393-0 x 10-5
`377-0 x 10-5
`350-0 x 10-5
`370-0 x 10-5
`405*0 x 10-5
`347*0 x 10-5
`370.0 x 10-5
`
`k1 +k2
`I In
`t
`
`K from
`equations
`k2 calculated
`Ka
`of Edgar
`from equation
`k,
`k2
`and Shiver
`Ka-(K+1)x
`PH
`(3)
`19*6 x 10-5
`1-37
`73-1
`52*3 x 10-5
`0-7 x 10-5
`53*0 x 10-5
`2 00
`72*8 x 10-5
`62-3
`71-4 x 10-5
`1*17 x 10-b
`74 0 x 10-6
`2-83
`7-85 x 10-5
`22-6 x 1o-5
`29-8
`242-0 x 10-5
`234-0 x 10-5
`311*0 x 10-5
`3-77
`5*673
`47-0 x 10-5
`326-0 x 10-5
`373-0 x 10-5
`1*335
`207-0 x 10-6
`4-64
`133-0 x 10-5
`311-0 x 10-5
`178-0 x 10-5
`0*086
`129'0 x 10-5
`5-63
`55-3 x 10-5
`76-5 x 10-5
`52*5 x 10-5
`18*2 x 10-5
`6-86
`44-2 x 10-5
`0-614
`27-2 x 10-r
`71-4 x 10-5
`46*1 x 10-5
`15*9 x 10-5
`28*6 x 10-6
`7-62
`0-613
`17-5 x 10-r
`8-49
`0*613
`1B 5 x 10-5
`23-9 x 10-i
`38*5 x 10-6
`14'6 x 10-r
`15*5 x 10-5
`25*7 x 10-i
`9.54
`15-8 x 10-5
`0-613
`41-5 x 10-5
`It has been possible to evolve equations based on equation (2) but involving
`three empirical constants which define with a fair degree of accuracy the
`relations of k, and k2 to pH within the range 2 to 10. They do not cover the
`changes in k, and k2 in strongly acid and alkaline solutions.
`The equations are
`A'k'[C'+[H+]]
`A"k" [C'+[H+]]
`(3),
`k -
`k
`[k"+[H+]][C+[H+]]
`[k'+[H+]] [C +[H+]] '
`2
`where A' = 3*68 x 10-3, C' = 0O8 x 10-6 and C= 1P9 x 10-5, A" = 2-25 x 10-3.
`It will be observed that, since K = k2/kl, K becomes
`0*6125 k" [[k' +[H+]]
`This is identical with equation (2).
`Equations of the sort developed above have little merit other than the
`approximate summary of a mass of data. In particular, one must be very
`cautious in assigning any material significance to the various empirical con-
`stants. One point cannot, however, be overlooked. The value of C is identical
`
`1
`
`008
`
`Harvest Trading Group - Ex. 1113
`
`

`
`928
`
`R. K. CANNAN AND A. SHORE
`
`with the dissociation constant of creatinine. It is difficult to see in what way
`the dissociation of creatinine can affect the intrinsic velocity of dehydration
`of creatine. A possible explanation might follow the assumption that there
`was involved in the reactions a tautomer having a constant similar to
`creatinine. C' might then be regarded as a second constant of this structure
`or as indicating the participation of yet another intermediary. A' and A"
`are merely the values which k1 and k2 would have were they determined only
`by the concentration of undissociated molecules of creatine and of creatinine
`respectively.
`
`2-5
`
`*
`0~~~~~~~~~~
`
`io_
`k1
`k2
`
`i
`
`10
`
`12
`
`1C 2
`4
`6
`8
`Fig. 2. Relation of velocityconstants to pHd
`0 Observed k2 from rate of dehydration of creatine.
`® Observed k2 from rate of hydration of creatinine.
`Observed values of kca are not inserted as they depart from their curve to the same extent
`as the corresponding values of k2t
`One final point of biological interest may be mentioned. Hahn and Meyer
`[1923] found that at 38f in a solution ofOh5 % creatine of pH7u01 there ap-
`peared an amount of creatinine corresponding to1h32 % of the total creatine
`in 24 hours. They calculate that the daily excretion of creatinine in the urine
`of an adult man corresponds to 1-33 % of the total creatine of his body. They
`suggest, therefore, that it is unnecessary to seek beyond the spontaneous
`dehydration of creatine for the origin of the creatinine of the urine. It follows
`from this view that the output of creatinine in the urine is governed only by
`the active mass of creatine in the muscles, by the temperature and by the
`PH of the muscle. Hahn and his associates have made an attractive case for
`this simple hypothesis. The data of the present paper give a general con-
`firmation of the above calculation. The velocity constant of the dehydration
`
`009
`
`Harvest Trading Group - Ex. 1113
`
`

`
`CREATINE-CREATININE EQUILIBRIUM
`
`929
`
`of creatine at PH 7*2 and 300 is 23 x 10-5. Applying the temperature correction
`of Edgar and Wakefield we arrive at a value of 43 x 10-5 at 380. This corre-
`sponds to the dehydration of 1*03 % of the active mass of creatine in 24 hours.
`This figure-somewhat below that of Hahn and Meyer-is sufficient to account
`for the daily otutput of creatinine provided the active mass of creatine in
`living muscle is as great as 0.5 %. Evidence continues to accumulate, how-
`ever, that this is an exaggerated value. It seems probable that only a small
`proportion of the total creatine which can be extracted from muscle by
`chemical means is in the free state in the living tissue. Unless the improbable
`assumption be made that combined creatine suffers dehydration as readily as
`when in the free state the argument of Hahn and Meyer cannot be sustained.
`It could then only be upheld were the demonstration made that the factors
`which have been shown to retard the velocity on the alkaline side of PH 3
`were partially suppressed in living muscle. It would be necessary for the
`apparent constant C to be diminished or the constant C' to be increased to an
`extent corresponding to the ratio of free creatine to total creatine in muscle.
`Upon this possibility there is no evidence.
`
`SUMMARY.
`1. Determination has been made of the apparent dissociation constant
`(uncorrected for. activity) of creatinine at 15°, 25° and 30° and of the first
`dissociation constant (uncorrected) of creatine at 170, 250 and 300.
`2. The velocity constants of the reversible system creatine-creatinine have
`been determined at 300 over the PH range 2 to 10, and have been related to
`[H+], the dissociation constants of the reactants and certain empirical con-
`stants.
`
`REFERENCES.
`Bjerrum (1923). Z. phy8ikal. Chem. 104, 147.
`Bronsted (1923). Rec. trav. chim. 42, 718.
`Cannan and Knight (1927). Biochem. J. 21, 1384.
`Clark (1922). The determination of hydrogen ions (Baltimore).
`Eadie and Hunter (1926). J. Biol. Chem. 67, 234.
`Edgar and Hinegardner (1923). J. Biol. Chem. 56, 881.
`Edgar and Shiver (1925). J. Amer. Chem. Soc. 47, 1179.
`Edgar and Wakefield (1923). J. Amer. Chem. Soc. 45, 2242.
`Hahn and Barkan (1920). Z. Biol. 72, 25, 305.
`Hahn and Fasold (1925). Z. Biol. 82, 473.
`Hahn and Meyer (1923). Z. Biol. 78, 91.
`Hunter (1928). Creatine and creatinine. (Monographs on biochemistry. London).
`McNally (1926). J. Amer. Chem. Soc. 48, 1003.
`Simms (1926). J. Amer. Chem. Soc. 48, 1239.
`Wood (1903). J. Chem. Soc. 83, 568.
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`010
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`Harvest Trading Group - Ex. 1113