'rARI,m DISSOCIATION CONbTANl'S AND
I
O r SEVERAL
~~IBONUCLl3OTIDmR
RET~ATET) COMPOUNDS AT 25"
Prcvious work is given in parentlieses. Compound ad en in^
A i PQI X 101 (Ni-H) 2.48 4.21 3.51 4.16 4.96 4.15 7.00 4.14 7.90 4.10 11.2 4.09 13.6 4.09 15.8 4.07 (4.1)" (4.15)b (4.12)'
ULTRASONIC DETTCRMINATTON OF REACTION RATES IN MAGNESIUM STJTJ'hTE ANT) MANCAXOUS SULFATE SOLUTIONS BY M. SUIZYANARAYANh Department of Physws, Nizam College, Hyderabad, I n d i a Reca'vcd Julu 18, 1961
The rate constants of the dissociation of magnesium sulfate and manganous siilfntc in aqueous solutions havc been determined by ultrasonic 9.88 (9.80) * methods by Bies' and by Kor and Verma,2 respec(9.75)O tively, using the theory of Bies (ref. 1). Thc theory of Bies, however, is based on the work of hIanes3 RP 6.32 6.60 in liquids and is extended to include electrolyte 8.02 6.50 solutions. Recently Tabuchi4 has worked out 9.85 6.60 14.0 6.34 a more detailed theory (applicable to electrolyte 17.2 6.31 solutions also) and has shown that the work of 19.8 6.46 Manes is only an approximation. Consequently, ADP 6.48 4.23 8.36 7.00 the rate constants obtained using the theory of 8.76 4.14 10.4 6.89 9.22 4.10 12.5 6.88 Bics also will be approximate. Employing tlhesame 13.0 4.16 17.6 6.84 experimental data, the author has, therefore, de15.9 4.08 21.6 6.80 termined the rate constants of the dissociation of 18 3 4.09 24.7 6.80 magnesium and manganous sulfate in aqueous (3.95)C (6.26)' (3.99)d (6.a~)~ Eolutions with the help of Tabuchi's theory. (3.9)s (6.65)" Because of absence of data Bies as well as Kor (6 68)' and Verma have taken the mean stoichiometric (6.3)' Adenosine 11.3 3.52 activity coefficients (J,tS) of zinc sulfate solutions 13.8 8.61 recorded by Harned and Owen6 as equivalent 15.9 3.47 to the mean rational activity coefficients of mag(3.6)" nesium and manganous sulfate solutions. The ( 3 . 63)' ( 3 . ii5)d mean rational activity coefficient (&) and the Rcsulta of A. G. Opton, J . Chem. Soc., 1713 (1!336), M dcgrcc of dissociation (8) of these solutions are not givon. b Results of 11. F. W. Taylor, J. Chcm. SOC.,765 calculated by the author, without any reference to 1'1948). .u = 0.001-0.007M . IIO added salt. Results of R. A. Albcrty, -R.- hr. smitd and 11. M. ~ o c l r lJ. Biol. zinc sulfate solutions, by the usual method of ive approximations. Since dilute solutions Chsrn., 103,425 (1951), = 0.15 M iXaC1. ResuIta of A. E. Martell and G. Srhwarzmhnch. Helv. Chim. Acta, 30, are involved the author has used the approximate 053 ( I H G ) , p = 0.1 M KCI. Results of N. C. hlelcliior, equation of Tabuchi's theory, namely, Twmax = J . Hid. Chon., 208, 616 (19*54), fi = 0.22 ill (C~ITG)IN+. Rrsults of Smith arid Alherty, rot. 10, fi = 0.2 M (n- 1. This approximation is justified by the earlier Ca€I,),N+. 0 Itraults of R. M. Rock, et al., Arch. B i o r h e m . work by the authors6 Riophys., 62, 253 (1956), fi = 0.10 M h'trc1. Bies and TCor arid Verma have found the rate constants to be invariant with concentration. sine and inosine. pH t,itration data in the present IIowever, employing t,hc proper values of the activcocfficicnt$sand using either of the theoricg it study indicate the proh:hlc dissociation of a proton ity is scen from Table I that the rate constant8 point with p K ca. 12 from RP, ADP arid adenosine; to a decrease with incrcafc of concentration. 'I'he however, glass clectrode mcnsiirements in lliis PIT decrease of lib, t,he rate constant for the association rcgioii are quite unrclinblc. It would be of iiiterest of inns, with the increase of concentration and honce t>o study this dissociation furthor ushg niethods of ionic strerigth is in conformity wilh the theory capable of yieldirig quant,itnativ(:results in this high PS explained by L~itllcr.' pII region. Tho molal volume changes, (AV)2, involved in The increased acidities of adenosine and RP rela- these reactions are calculated for these solutions tive to that of adenine or I-12P04- ( p K = 7.20) in- by using the formula given by Ries, where the condicate that the ribose acts as an electron withdraw- centration has to be expressed in moles per e ~ . ing group. It is interesting to note in this connec- The values obtained for magnesium and manganous tion that pK1 is the same in ADP as it is in adenine (1) D.Ries, J . Chdm. Phye., 23, 428 (1955). (2) S. K. Kor and G. 8. Verma, ibid., 29, 9 (1958). and p K 2 is appreciably larger in ADP than it is in (3) M.Manes, ibid.. 41, 428 (1953). RP. The increased pKp value probably is due t o (4) D.Tabuchi. ibid., 26, 993 (1957). thc increascd negative charge on the ADP relative ( 5 ) H.R. Harned and B. B. Owen. "Physical Chemistry of Electo that on RP. This is further substantiated by trolyte Solutions," Reinhold Publ. Corp., New Pork, N. Y., 1950, 190,426. the increasing p& values observed in the series p. (6) M.Kriahnamurthl and M. Suryanarayana, J . Phys. SOC.Japan, adenosine mono-, di-, tri- and tetra-phosphate.12 16, 2318 (1960). 14.0 17.0 19.8
9.66 Q.6Q
I
\ - - - - , ,
(12) R. M.Smith and R. A. Alberty, J. Pliys. Chem., 60,180 (1956).
(7) K. Laidler. "Chemical Kinetics," McGraw-Hill Book Co., New Pork, N. Y., 1950. p. 132.
Noms RATECON8ThNTR
FOR IONTC n1RSOCIATTflN IN
‘I‘A~LF, I AQTTEOI:S 8OLUTIONS
OW
hf A C N E S I U M
MhNoaNous
SULI”ATE! A N D
SULFATE
(250)
Concn. (mole/l.)
Salt
Magnesium sulfate
diesociation, 6
0.003 .005
0.851
,008
.776 .760 .734 .707 .go2 ,851 .779 .717 ,662
.010 .014 ,020
Manganoussulfate
Degree of
.001 ,002 .005 ,010 .020
Mmn rntional nctivity cocfficient
.809
-Tabuchi’s
f*
0.658 .594 .542 .515 .473 .430 .770 ,705 ,604 ,522 .439
5.74 X 10’ 5.52 5.01 4.43 4.35 4.30 12.6 x 106 11.57 10.32 9.21 8.24
sulfate solutions are of the order of 7.7 and 120 (cc./mole)2, respectively.
ON THE REACTIVITY OF HYDROGEN ATOMS IN AQUEOUS SOLUTIONS BY JOSEPH RABANI Department of Physical Chemiatru. The Hebrew unite rail^. Jerwalem, 1 8 T d
Receited Julu 19. 1961
Hart’ measured the constants ratio kl/kz for the reactions where HC02H denotes both formic acid H + +HOs + HCOzH +Hz + COzH 0 2
13
(1) (2)
and formate ion. Using various concentrations of formic acid and oxygen, it was shown that at low formic acid concentrations the ratio kl/k2 is about 500 a t p H -3. This value was used for the estimation of rate constants218a t other pH values. However, a t the high formic acid concentration (1 iM)this ratio is about 5000 (pH 1.7-1.8).’ This discrepancy is due to the different reactivities of IICOzH and I-IC02- toward hydrogen atoms.4 Using various concentration ratios of formic acid and ferricyanide, the yield of the hydrogen produced by the action of X-rays was measured a t the p1-I range of 1-3.4 In this system, formate and ferricyanide compete for the hydrogen atoms according to reactions 2 and 3. H
+ Fe(CN)s*- +H + + Fe( CN)64-
theoryki,, I./molo-sec.
krl BCC.-’
(3)
The ratio k3/k2 is pH dependent. At pH 1.3 this ratio is about 3500, falling to 800 a t pH 2.05 and to 300 a t p H 2.55. These results were obtained in solutions of constant (0.047 M ) formic acid concentration, for direct comparison with Hart’s experiments. These results indicate that in the ferricyanideformic acid system, only ferricyanide and formate ion (and not the undissociated acid itself) react with H atoms, a t the pH range investigated. Reaction 2 has to be replaced by (2’). (1) E. J. Hart, J . A m . Chrm. Soc., 7 6 , 4312 (1954). (2) J. H . Baxendale and D. 11. Smithies, 2. phyaik. Chcm. (Frankfurt). 7, 242 (1956). (8) P. Ricsz and E . J. Hart, J . Phye. Chcm., 6 3 , 858 (1959). (4) J. Rabani. P1i.D. Thesis, The Hebrew University, Jerusalem, Israel, 1961.
9.11 8.76 7.95 7.03 6.90 6.87
x
107
2.52 x 100 2.31 2.07 1.84 1.65
H
--
B i d thcory-
ki. sec. -1
kh, I./molesec.
6.37 X 105 6.47 6.19 5.64 5.82 6.16
10.1 X lo7 10.3 9.8 9.0 9.2 9.8
13.2 x 106 12.63 12.25 12.06 11.72
+ HCOs- +H1 + COz-
2.64 x 109 2.53 2.45 2.41 2.35 (2’)
Kinetic evidence indicates that the oxidation of formic acid in solution by C12,6 permanganate,s some other inorganic ions6 and SOr- radical’ mainly proceeds by the oxidation of the formate ion. The values of k3/k2’ were found to be 12.5 a t pH 1.3, 15.4 a t pH 2.05 and 17.9 a t pH 2.55. Thus, k3/k2’ does not depend on t,he pH, within experimental errors, in agreement with the present interpretation. In Table I. our results for the ferricvanide-formate system are compared with those for the oxygen-formate system.’ TABLEI THEREACTIVITY OF H ATOMSIN FORMIC Acm SOLUTIONS kdktl
ki/kP
kdki
1.71 3950 1.83 5600 2.32 1875 2.35 945 3.12 526 a Computed values.‘
1700 1300 430 385 81
2.3 4.3 4.4 2.4 6.5
PH
Taking into account the statistical weight of these data, a mean value of kl/k3 = 5 is obtained. Since k3/k2 is 3500 in 0.1 N HzS04, it follows that kl/k2 = 17,500 a t the same pH. Using the recenb + X lo6 results of Riesz and Hart3 for k ~ + F e ~ (4.8 1. mole-‘ sec.-l) a t pH 2.1, the following corrected values for the reactivity of H atoms are obtained (Table 11). These “absolute” values are based on gaseous phase data. The numerical values in column 2 of this table represent relative rate constants with respect to 0 2 : k~ + s/kH + ormultiplied by lo4. The values taken from reference 4 were obtained from measurements of hydrogen yield in the ferricyanide-organic solute systems. Much care should be taken in considering the reactivity of hydrogen atoms in aqueous irradiated ( 5 ) J. Thamaen, Acto Chem. Scand., 7 , 682 (1953). (6) J. Halpern and 9. M. Taylor, Diecussions Faradou Soc., 174 (1960). (7) E. J. Hart, J . A m . Chem. Soc., 83, 567 (1961). (8) W. G. Rothsohild and A. 0. Allen, Radiolion Reeearch, 8, 101 (1958). (9) J. H. Baxendale and G. Hughes, 2. phyaik. Chem. (Frankfurt), 14, 306 (1958).
