1900
- 1
L 1-
-
Vol. 65
hTOTES
30
-31
1
I
-
0
-L
10
20
30
en-'. E!g. 1.
rolved in bondiiiq, the equilibrium constant is of different orrlt r from that of IL-hydrocarbon complexes. 111thc present case the aza-complexes and hydrocarbon complexes approach each other in stability more closely. It seems that both n-7r and T - A interaction play their part in the formation of aza-complexes, the contribution of A-ir interaction being more i r ~the present case than for 11complexes. For the determination of equilibrium constants, we have used tbe modified Benesi-Hildebrand relation, i.e.
have been determined at 2 5 O on several occasions, there is very little information avltilable about the enthalpy changes of these reactions. Gallagher and King' recently have reported enthalpy results determined calorimetrically for the case with chloride ions. We present here similar results for the reactions with chloride, bromide and iodide ions, all at 2 5 O . The results were obtained by measuring the heat effects produced by the dilution of small amounts of potassium halide solution with large volumes of acidified mercuric perchlorate solutions. The details of the method have been described previously.a The ionic strengths of the various final solutions were between 0.10 and 0.14. In al! cases the concentration of the species HgXf was between twenty and a hundred times greater than that of the species HgX2, and the concentrations of the higher complexes were negligibly small.
extinction coefficients of complex and ncreptor sut)strate, reap. 0ptic:tl density oi niisture (A) E = - _,; . OpKl\!:tl path leiigth (1) X added initin1 concn. of arccptrt: (C,O) CD= conreiitration of donor
Experimental Apparatus -The calorimeter has been described already .* I n each elperiment 10 ml. of potassium halide solution was diluted with 990 ml. of mercuric perchlorate solution The calorimeter was calibrated again on two separate occasions during the course of this work by measuring the heats of the one-hundredfold dilution of two different sodium chloride solutions. The average values obtained for these heats of dilution a t 25" (with the corresponding literaturea values shown in brackets) were 302 (297) and 316 (326) cat. mole-' of sodium chloride for solutions with initial concentrations 2 50 and 2.73 molal. The calorimeter was immersed in a thermostat a t 2 5 0.01'. Materials.-A.R. Chemicals were used throughout the uork. All stock solutions were rrade up by weight. -4 stork solution of mercuric perchlorate was prepared by dissolving a known weight of mercuric oxide in an excess of 60% perchloric acid. The pH of the solution wa8 measured and the concentration of mercuric ion wss checked by analysis. The concentrations of mercuric ion in the solutions used for the dilution of the potassium halide solutions ranged from 0.032 to 0 050 molal.
The plot of ] / ( I < - &) zzgaiiist 1 j C ~givos x straight line, the intercept of which on the I / ( E - E,) axis gives 1/(& - E,) and slope of which gives l / K ( E , - E=). Recently Xash5 criticized the Benesi-Hildebrand relation on the ground that an auxlliary quantit,y, the molar absorptivity of the comp!es, need h e determined before the object of primary iritercst, the equilibrium mnst.nnt, can be determincd. We have avoided this diiEculty by dcterrnining .K directly from tlie intercept on thc I/CD nsis. Xash's eyuation5 can, however, be derived very easily by simple algcbrai r transformatioil of the nbove e q w t ion. T1larlk.c arc d u c to Professor S. Basu for helpful cliscassioiis.
Results The results of the dilution experiments are \hewn in Table I. The concentrations of the ions in each solution IT ere calculated from the total concentrations of mercuric and halide ions gwen in Tablr 1 and the nworintion constants of the cornplex ions The assoriation constants used were those drtermincd by Marcus4 at 2 5 O and ionic strrngth 0 50, and nere roorrccted to the particular ionic strength\ of the yoliitions by the equation of Davies.5 Thc rrlntion betwcwi the heat of associntion of reaction 1 at zero ionic strength and at any other ionic strength for which the association constant is K' is g v c n I n 7 the formula6
whcre E,, E,
-
-
(5) C. P. \id!, J . Pitun. Cham., 64, 9 3 (1960). I -
The calculrited values of A H 0 for the thrce nrFociiI(1) P. K. Gsllagher a n d E. L. King. J. A m . Cham. S o c . . 82, 3510 (1960).
J., P h w (2) n. W. Anderson, G . N. Malcolm and 1%. P;. P a r t ~ ~ n Chrm., 64, 494 (1960). (3) "Selected Values of Chemical Thermodynnrnic I'ropcrtirs." Piatl. Bur. Standarch Circ. No. 500, 1952. :4) Y. Marcus, Acto Chem. SronJ., 11, 599 (195;) ( 5 ) C. u'. Davies, J. Chem. Soc., 2093 (1'338\. ( 6 ) J. ?VI. Austin, R. A . &%:heson a n d H. N. Parton, "l'i~e Structure of E!ectrolytio Solutions," edited by \$'. J. IIainer, J i ~ l i uWiley nnd Sonu, New York, N. Y.,1059.
NOTES
Oct., 1961 TABLE I THEHEAT(Q)OF THE REACTION diluted with KX(m) --+KX(O.01 m) Hg(ClO&(m') 711
1 .oo
0.50
0.25
Ionio
Moles of IC1 x 1 0 %
x 10.11 10.00 9.97 4.74 4.72 4 69 2.48 2 48 2.48
=
.047
.046
=
10.88 10.79 10.80 4.74 4.73 4.75 2.28 2.29 2.25
0.046
1 00
10.90 11.10 11.0
0.042
0.50
5.41
.050
0 23
5.28 5.42 2.70
.050
0.50
I),
25
atrength of final Boln.
c1
0.047
X 1 .OO
Q,cal.
m'
-57.07 -57.07 -57.09 -29.3 -29.2 -28.1 -113.4 -13.0 -13.0
0.13
-96.6 -95.9 -96.0 -48.5 -48.1 -48.4 -25.0 -24.6 -24.0
0.13
-184.7 -181.1 177.8 -89.7 -89.8 -92.5 -49.2 -52.3 -55.1 -10.7 -18.4 -1!1.2
0.12
.13
.13
Br
.032
.032
.09
.09
x-I
0.10
-
3 i3 L' 74 0 04 0 (32
.0'12
0.92
.I5
.15
.I2
TABLE I1
FENCTIONS FOR THE REACTIONS Hg+? 4-X - = HgX+ at 25'
'rlIICRMODYNANIIC
CI Br ;
.3G'.
kcal. r n o l t - ~ ~ 1
-10.0 -13.3 -18.4
AHQ
AS0 expt..
kcal. m6le-1
e.u. mole-1
-- 4 . 8 f 0 . 5 -10.6 f .5 -17.6 f . 5
17 f 3 9 i3 3 5 3
whereas their values are required at ionic strengths 0.1 and zero. In the absence of more reliable information the Davies equation was used to extrapo!ate the association constant values from ionic strength 0.5 to zero. The reliability of this equation a t ionic strengths greater than 0.1 is not known. Nevertheless it is of interest that Vanderzee and Dawson7 have used an equation very similar to the Dsvies equation to fit association constant values for CdCl+ measured a t ionic strengths 3, 2, 1 and 0.5, and obtained an extrapolated value a t zero ionic strength in good agreement with an independent determination hy Harned and Fitzgerald.8 In all the final solutions the halide ion is present, predominantly as HgX+ so that, the values of A H do not depend to any great extent on the values of the association constants. The correction of the enthalpy values to zero ionic strength involves the use of the Davies equation. The magnitude of this correction is never more than 0.4 kcal. mole-'. It is probable therefore that the uncertainty in the standard enthalpy values is within f 0.5 kcal. mole-'. The standard entropies of association will be in error by =k 2 e.u. mole-' from this cause. If the uncertainty in the free energy values arising from the use of the Davies equation i s taken into account, the error in the stand;trd entropy figures rises to as much as f 3 e.u. mole-'. Gallagher and King' obtained a value for thc enthalpy of association of HgC1+ of -5.9 kcal. mole-l at ionic strength 0.5. Correction of t h k figure to zero ionic strength by means of equation 2 gives a value of -5.4 kcal, mole-'. The agrc.tbment between the two result,s for AH*(IlgCI+) i: not close, but is within the combined experimental rincertaint ies. Entropies of Association.-C'omup:~risoii ( J f t,hc results in T&le 11 with those for the correspontliri:: reactions with cadmium(I1) and lead(I1) ions9 rcve:ile that the entropy c,hangcs for the association of eac,h cation with a given ha!itle :inion arc: f,li(. same within t.he cxperiment,nl responding cu t,halpics of assocint i c ~ i i :we markedly different. ]:or exnmp!e, for t,he rcnctions invol..iiig the formition of HgBr+, Cdl3r-+:ind PbRri, the entropies of association are all ai)o\it 9 e.u. moIc--l but the enthalpies of associ:tt ion are -. 10.6, --0.3 aid f0.3 kcal. mole-l, rcspcctivcly. Thew results arc consistent with t,he that the species HgX+, CdX+ and PbX+ are cornplcx ions with direct contact bctwcen the cation and the mion, RO that their formation involvcs R i.h:lriFP in hydi~3tion structure of the ions. For ~1 qi;-i.n anion thi: strengths of the various cation-anion horids will influence the enthalpies of the association re:i,ctions but. necd have iittle effect on thc mtropica of association if these are Inrgely detwmmcd by the change of hydration of t,he inns CJR association. George'O has shown hhat t,he ent,rripicx 14 a aumher of cation-anion association rescf-iorls C R T ~tic r u m st.1itt.d by the equ at'ion (1
t,ion reactions i rivolving the inercury(I1)-halide c+omplexions are shown in Table IT. The st,and:u.d free energice of association ealculatcd from the msociation constants a t zero ionic strength, and the standa'rd entropies of association calculated from t'lw free energies and enthalpies are shcwn in t,hc same table.
S
1901
A S 0 calc.. e.u. mole-'
17 11 4
Discussion Errors.--Th;is method of determining enthalpies of association of complex ions depends for its ,-:i.iccess on a knowledge of the corresponding association constants at, the appropriat,e i o r k ,strengths. In the prwent case the msociatior! t-rmstants w e known only at ionic strength 0.5
v i m 7
( i ) C. E. Yaridersec and €I. J. D s w e a u , J . Am. Chsm. Sot 7 6 . 5659 11953). (8) 11. E. Hsrned and M. E. Fitrperdd. i b d , 68, 2624 (j'W3. (9) G. 11. Nancullaa, QUUT~. Reca. (London:. 1 4 , 402 (1WC). (IO) J. D. U. George.. J . Am. Chem. s'cc.. 6 2 . 5530 (1959'
NOTES
1902 ilS"
:=
A S (hydration of anion)
+ constant
(3)
The entropy results for the mercury(I1)-, cadmium(11)- and lead(I1)-halide systems conform to this pattern with constants -9, -10 and -9, respectively. The implication of these results according to George is that these association reactions involve the loss of water of hydration. An equation for the entropy of association has been derived by Matheson6 based on an entropy cycle involving the replacement of one water molecule in the hydration shell of the cation by the anion. The form of the equation for water a t 25' is
where Z1 m d Z 2 are the charges on the central cation and the complex ion, respectively, and T + is the crystal ___ __ radius of the cation in hgstroms. SOH~Oand SOX- are the standard partial molar entropies of the water molecule and t,he anion in aqueous solu.tion.ll When the value pf r+ for equathe mercury(I1) cation is taken as 1.10 tion 4 leads to the entropy values shown in the fifth column of Table 11. Alt,hough the equation involves several approximations there is good agreement between calculat'ed and observed values. Siniilar agreement was found when the entropies of association calculated by equa.tion 4 were compared with the observed values for the I : 1 coniplex .halides of cadmium and lead.6 (11) R. E. Powell and W. M. Latimer, J . Chem. Phys.. 19, 1139
(1951).
(12) It. -4. Robinson and R. 11. Stokes, "Electrolyte Solutions," Butterworths, London. 1959.
PROTON MAGNETIC RESONANCE OF GLYCYLGrLYCINATE, GLYCINEAMIDE, AND T H E I R METAL COMPLEXES BY
NORMAN
c. LI,
I & h Y JOHNSON AND JAMES SHOOLERY
Varian A8socaates, l'nlo Alto, Calif.,and Duqzlesne Universitz, Pitlsburgh. P a . aecewed April IS, I961
Several papers' , z have appeared recently dealing with proton resonance studies of dipeptides. Takeda and Jardetzky' found that in glyoylglycine there are two peaks from the two CH2 groups, indicating tha,t they are non-equivalent. It, would be of interest t'o cont~inuesuch studies of dipeptides and determine the effect of chelating metal ions 011 the chemical shifts of these two CH2 groups. A Varian Associates Model A-60 NMR spectrometer which operates a t 60 Mc. was used. Solutions containing 1 .Z M concentrations of each coiistitueiit in 9Y.Sr)& TIzO were placed iii 3 mm. i d . , 4 rnm. o.d. precision Pyrex tubes and the tubes were filled to a depth of about 40 mm. Pure tetramethylsilane in a 5 nun. 0.d. annular cell was used as an external reference compound. The accuracy in measuring the peak positions is estimated to be =tO.Ol p.p.m. The resonance signal of the protons being measured occurs a t lower field than that of the reference, that is, H < H,, and the ( 1 ) M. Takeda and 0. Jardetzky, J . Chem. Phya., 96, 1346 (1957). (2) F. A . Borey and G. V. D. Tiers, J . Am. Chem. Soe., 81, 2870 (1959).
Vol. 65
dowiifield chemical shift is taken to be positive, in accord with the definition 6
=
1O6(H, - H ) / H ,
where Hr and H are the resonant fields for the reference substance and the sample, respectively. PROTON
TABLE I DUE TO METALIONS
CHEMICAL SHIFTS
Volume susceptibility,
THE
PRESENCE OF
Sample
x6
CHz chemical shifts, p.p.m.. (values in parentheses corrected to xv for GG-)
Sodium glycylglycinate, GGZnClz GGhfgClz GGGlycineamide, GA ZnClz GA
0.756 .785 .768 .749 ,766
4 . 2 5 ( 4 . 2 5 ) 3.84(3.84) 4 , 4 2 ( 4 . 3 6 ) 4.20(4.14) 4.35(4.33) 4.00(3.98) 3.80(3.82) 4.20(4.18)
+ + +
Since an external reference is used, bulk-diamagnetic susceptibility corrections must be applied.3 For a cylinder of length large compared with the radius, the equation applicable here is &om.
=
Bobsd.
+ 2 ~ / 3 ( x v-
Xv.r)
where xv and x ~ are , ~the volume magnet.ic susceptibilities of the sample and reference, respectively. The volume susceptibilit'y of each of the solutions was determined immediately after the n.m.r. spectrum was taken, using the same spectrometer, following the method of Reilly, et u Z . , ~ based on the use of the non-rotating concentric The sample whose susceptibility was to be determined was placed in the inner cylindrical space and a substance with a single sharp line (in this case tetramethylsilane) was placed in the annulus. The resonance from the material in the annulus displays two maxima whose separat,ion ( n c.P.s.) is a linear function of the volume susceptibility of the liquid in the inner cell. The values of n for three pure liquids: chloroform, carbon tetrachloride, methanol (at room temperature, -xv = 0.735,6 0.684,3 0.530,' respectively) were determined and a linear plot of n vs. xv was obtained for the three liquids. From this graph, and from the values of n obtained for the solutions, the magi l o t i r susceptibilities of the latter were obtained. Tiiese are included in Table I. The use of an 1i.m.r. spectromet'er for determining magnetic susceptibilities consumes much less time than the use of the c*onventional Gouy method, and is recomnicnded for routine susc:ept,ibilitymeasurements. In Table I, the two lines in GG- are due t o tho two iioii-equivalent CH2 groups and the dowiifield chemical shifts are increased by 0.11 and 0.:30 p.p.ni., respectively, in the prcscnce of 1.2 M ZnCln. When ZnClz is added to a solut,ion of sodium glycylglycinatc iii D20, a chelate is formed, and there is either a reduction of electron deiisit'y (3) J. A. Pople, W. G . Schneider and H. J. Bernstein. "High Resolution Nuclear Magnetic Resonance," McGraw-Hill Bonk Co., New York, N. Y., 1959. (4) C. R. Reilly, H. M. McConnell and R. C. Meisenheimer, Phys. Res., 98, 264A (1955). ( 5 ) Rilmad Glass Company. (6) A. A. Rothner-By and R. E. Glick, J. Chem. f'hy8., 56, 1652 (10.57).
( 7 ) S . Broersina, ibid., 17, 873 (1949).
