April, 1951
POLAROGRAPHIC
REDUCTION OF C O P P E R ETHYLENEDIAMINE TETRAACETATE
330
THE POLAROGRAPHIC REDUCTION OF COPPER ETHYLENEDIAMINE TETRAACETATE BY E(. BRILAND P. KRUMHOLZ Eesearch Laboratory of Orquima S.A., Sdo Paulo, Brasil Received Noaember 24, 1853
A functional dcpendence of the half-wave potential of copper ethylenediamine tctraacctate (Hnta) upon the pH, i n well buffered solutions, is discussed. The formation of multiple waves in unbuffered solutions of copper Enta, containing small amounts of free hydrogen or metal ions capable of forming stable complexes with Enta has been studied. An approximate treatment of this effect is given based on the Ilkovic theory and the assumption of depletion of hydrogen or metal ions from the solution-mercury interface. The interference of kinetic effects is stated.
It is known that the copper ethylenediaminetetraacetat'e (Enta) complex produces a well defined polarographic wave. l Its half-wave potential in buffered solutions is shifted toward more negative values with increasing pH.'" Addition of small amounts of acids or metal ions, forming stable complexes with Enta, produces in unbuffered solutions of the copper-Enta complex a peculiar wave splitting.2 Similar double or multiple waves have up to now been chiefly observed in organic systems.a I n the following we report on this particular behavior of the copper-Enta system and its possible explanation. Experimental The apparatus used was described in a previous communication,2 as was also the preparat,ion of Enta, copper, lead, cadmium, nickel and zinc solutions. The calcium nitrate solution was standardized according t80 the oxalate-permanganate method. The lant.hanum nit.rat,e was prepared from a pure oxide (purity 99%) calcinated at 950". The dropping mercury electrode had m'lat'/B = 2.53 ing.Vr sec.'/e as determined in a solution 0.1 molar in potassium nitrate, 5 X 10-6% in methyl orange, oxygen free, at 20.0"in short circuit. Current measurements were made in general in 10-mv. potential int,ervals. The pH of the solut,ions was maintained by the addition of potassium hydroxide or nit,ric acid with potentiometric control of the pH using a glass electrode mounted directly in the polarographic vessel. The pH of all solut#ionscould thus be checked during the actual measurement. The reproducibility of the pH control in buffered solutions as well as in the unbuffered solutions in the H range 3.5 to 4.54 was about f 0 . 0 2 pH unit. In unbuzred solutions in the pII range 5.5 to 6.5 the stability of the pH readings was about 0.1 pH unit. All solutions were brought to cquilibrium with respect to the exchange reaction (see ref. 2 and 4) CuYMecz,"+ = Me(z,Y(4-")Cu++
+
the concentration of the supporting electrolyte. At constant ionic strength, it shifts toward more positive values with increasing buffer concentration (see Fig. 1). At constant buffer concentration it shifts toward more negative values with increasing ionic strength. Beyond the dotted line in Fig. 1, both the buffer concentration and the ionic strength increase, and it is actually probable that the apparently constant value of the half wave potential is only a result of a counterbalancing effect. Figure 1 shows that a t low buffer concentrations the half-wave potential is almost independent of the concentration of copper-Enta. With increasing buffer concentration there appears an increasing shift of the half-wave potential toward more negative values as the copper Enta concentration is increased from 10-4 to loA3 molar.
1
0
~~~~
0.20 Molarity of the buffer. 0.10
-0.30
+
Experimental Results The half-wave potential of the copper-Enta complex a t a given pH depends strongly upon the concentration of the buffer used, as well as upon (1) (a) R. L. Pecsok, J . Chem. Educ., 29, 597 (1952); (b) H. Ackermann and G. Schwarzenbach, Helu. Chim. Acta, 36, 485 (1962); (0) W. Furness, P. Crawshaw and W. C. Davies, A n ~ Z y s l74, , 629 (1949); (d) J. Koryta and I. Kossler, CoZZeclion Czech. Chem. Communs., 16, 241 (1950); (e) P. Souchay and J. Faucherre, Anal. Chem. Acta, 8 , 252 (1949); ( f ) BUZZ. soc. c h i n . France, 7 2 2 (1949); (9) R . Pribil and B. Matyska, Chcm. L i s f y , 44, 305 (1950); (11) E. I. Onstott, J . A m . Chena. Xoc., 74, 3773 (1952). (2) K. Bril and P . Krumholz, THISJOVRNAL,67, 874 (1958). (3) (a) I. M. Kolthoff and J. J. Lingane, "Polarography," 2nd ed., Intersoienoe Publishers, Inc., Chapter XIV and XV. (b) M . von Stackelberg, "Polarographische Arbeitsmethoden," W. de Gruyter, 1960, p. 309, etc. (4) Actually in this p H range some buffering occurs due to t h e acid function of the copper Enta complex; see G . Schwareenbach and E. Freitag, HeEu. Chim. A c ~ 34, , 1492 (1951). We found by potentio= 1 . 2 X 108 a t 20' and 0.1 ionic strength. metric titration
0.025 0.050 0.075 Molarity of the buffer. Fig. 1.-The influence of the buffer concentration upon the half wave potential of the copper-Enta complex. Left of the dotted line, the ionic strength is maintained constant ( f i = 0.10) by suitable addition of potassium nitrate. Right of the dotted line the ionic strength, which is given by the buffer concentration, increases proportionally to the buffer concentration: A, CHICOOR CHsCOONa (pH 4.69); B, NaHaPOa Na2HPOI (pH 6.88); C, potasNa2C03(pH sium tetraborate 6pH 9.32); D, NaHCOg 10.08); t = 20.0 . (All solutions contained 5 X 10-6'% methyl orange.) 0
+
+ +
K. BRILAND P. KRUMHOLZ
340
pressor in the concentration range of 5 x to 15 X By further increasing this concentration the prewave is depressed and suffers a shift to more negative potentials. This effect is essentially the same for all suppressors used.
This shift reaches about 30 mv. for solutions where the total ionic strength is given by the concentration of the buffer. A further increase of the buffer concentration (and of the ionic strength) has no significant effect upon this shift. The wave slope
GO
TRATION
OF
0
2C 30 .2
CONCENTHE BUFFER FOR A 10-aM COPPER-ENTA SOLUTION AT u = 0.1, 20.0" CHsCOOH CHaCOONa 4.69
Buffer PH Total molar concn.
0.24" .20 .10 -05 .037b .025 .020 .012 .0062 a p = 0.12
KHePOd KaHPOd 6.88
HsBOa NaHpBOa 9.32
I
fa 40 .-
TABLE I THE
--
50 -d, .-
in buffered solutions of constant ionic strength depends also upon the nature and the concentration of the buffer, decreasing slowly as the buffer Concentration is increased, as shown in Table I. DEPENDENCE OF THE WAVESLOPE,0 , UPON
VOl. 58
5 u
.20 .-
__
10 .-
NaHCOa NazCOa 10.08
0
a
0.030 .031 .038 .039
0.035 0.030 0.032
*
0.040 0.042
0.036 0.037
0.035 .036 .036 .039 .043
= 0.15.
In unbuffered solutions of the copper-Enta complex containing very small amounts of free acid the polarographic wave shows a peculiar wave splitting illustrated in Figs. 2 and 3. Two distinct pre-waves appear clearly. The height of these waves, a t constant copper-Enta concentration is approximately proportional to the hydrogen ion concentration.
--
I
I _-
I
250 500 E us. S.C.E. in mv. Fig. 2.-Wave splitting in slightly acid unbuffered solutions of copper-Enta in 0.1 M potassium nitrate, copperEnta concentration 10-8 M , 5 X 10-6% methyl orange, t = 20.0": 1, copper-Enta wave at pH 7.1; 2, copper-Enta wave a t pH 3.99; 3, the calculated currentp-otential curve (for E1 the value -464 mv. was used).
..
I
0
In the absence of maximum suppressors two very prominent maxima appear in these solutions, as shown in Fig. 3. Methyl red, methyl orange, anaphthol, caffeine and gelatin, all can be used to suppress these maxima. For methyl red and methyl orange the polarographic wave is independent of the concentration of the maximum sup-
0
250
500 750 1000 E us. S.C.E. in mv. Fig. 4.-Wave splitting in a 0.4 mM copper-Enta solution, 0.2 m M in metal nitrate, 0.1 M in potassium nitrate, t = 20.0", pH 6.1 += 0.2: I, lead nitrate; 2, lanthanum nitrate; 3, cadmium nitrate; 4, zinc nitrate; 5, calcium nitrate; 6, no metal added. Abscissa scales for curves 2, 3, 4 5, 6, start at 100, 200, 300, 400 and 500 mv., respectively, to the right of curve 1. (All solutions contained 5 X 10-5% methyl orange.) (5) N o wave splitting could be observed upon addition of nickel salts.
April, 1954 0
POLAROGRAPHIC REDUCTION OF COPPERETHYLENEDIAMINE TETRAACETATE
341
other metals, and of about 60 mv. for the hydrogen induced wave.
Mmoles La+++/]. 0.2 0.4
300
400 500 E us. S.C.E. in mv. polarogram of a 4 X
600
M copperFig. 6.-Derivative Ent: solution containing 2 x lo-' 111 calcium nitrate a t 20.0 ; = 0.1 ( 0 ) ;p = 0.5 (0). 0
250
500
E us. S.C.E. in mv. Fig. B.-Wave splitting produced by various concentrations of lanthanum nitrate in a 0.4 mM coopper-Enta sohtion in 0.1 M potassium nitrate a t 20.0 , pH 6.2 0.1: 1, current-potential plots; 2, diffusion current (induced by the metal)-concentration plot.
+
Figure 5 shows the dependence of the height of the metal induced wave upon the concentration of the metal ion in the case of lanthanum a t pH 6.2 f 0.1. A very good proportionality is observed. The same was found for all other metals studied here.6 The height of the metal induced wave shows no marked dependance upon the concentration of the copper-Enta complex. It has a normal temperature coefficient as shown by the fact that, all other conditions being maintained constant, the relation between the height of the first wave and the total height of the wave is, within the experimental error, independent of the temperature in the range of 0" to about 50". I n the case of the metal induced wave as well as in that of the hydrogen prewave the total ionic strength exerts a pronounced effect upon the wave separation: the higher the total ionic concentration the poorer the wave separation. Thus in the case of calcium, the double wave formation apparently disappears, in 0.5 M potassium nitrate. It manifests itself only by broadening the derivative polarograms (see Fig. 6). As shown by this figure the half-wave potential of the second wave remained almost unaltered by the change in the ionic strength, whereas that of the first wave is shifted by about 40 mv. toward more negative potentials with increasing ionic strength. Values ranging up to 70 mv. (for lead) were sound for the (6) Actually the metal induced wave is proportional to the concen-
tration of the free metal ion. This, in general, is different from the MewR = Mewadded, because of the exchange equilibrium CuYy'4 -70Cu + +. Only in the case of lead, however, does this correction become important. I n all other cases studied here i t can be neglected, within the experimental error, when the wave height of the induced wave is measured from zero current line ut) (corrected for the residual current). This is due t o the f a c t t h a t the diffusion coefficients of the metal ions studied here, with the exception of lead, are almoat equal (see ref. 3% p. 52) to each other (see also ref.2).
+
+
+
Some experiments were performed using the streaming mercury electrode for comparlson (see ref. 2). The wave separation is in all cases poorer a t this electrode than a t the dropping mercury electrode. Thus, in the case of zinc even in 0.1 M potassium nitrate the wave splitting at the jet electrode manifests itself only on the derivative diagrams. Raising the temperature seems to improve the wave separation. Only lead gives a t the jet electrode a clean wave splitting, a good proportionality existing between the free lead concentration (see ref. 6) and the height of the induced wave.
Discussion
A. Well Buffered Solutions.-The diffusion current constant of the copper-Enta complex is only about 10% smaller than the diffusion current constant for copper in nitrate medium.' In acetate solutions the difference observed is of about 5% only. This indicates a two-electron electrode process which can be tentatively formulated as CuY-
+ Hg + 2e = CuHg + Y4-
(1)
Except in very alkaline solutions the Enta ion, Y4-, will react with hydrogen ions, free or furnished by proton donors as buffers or water, according to Y4- + nH+ = YH,(44-. . .KYH,, (2) Assuming as a first approximation the reduction to be reversible, the potential of the mercury electrode should be given a t 20" by
where F is an activity factor and the (X), indicates concentrations a t the electrode surface. If reaction 2 as well as the dissociation of the proton donors occurs instantaneously, then the concentration (Y4-)0 of the Enta ion a t the interface can be expressed as a function of the total analytical concentration of Enta a t the interface (Y*)o,the concentration of the hydrogen ions, and (7) See also ref. Xb, P. 487.
K. BRILAND P. KRUMHOLZ
342
of the acid dissociation constants of Enta* according to A
(YHn)o = (Y")o f (H+)o
(Y*)o =
(4)
0
where f(H+)o
=
1
+ Ki(H+)o+ KiK2(H+)02+ &K2&(H+)OS
+ ..
(4a)
By following a treatment similar to that proposed by Tomesg in the discussion of the polarographic reduction of mercury cyanide equation 3 can be put into the formlo E = El
+ 0.029 log f (H+)o- 0.029 log
zd
-a
(5)
Equation 5 presumes the validity of the Ilkovic equation for the diffusion of all components of the electrode process. Equation 5 predicts a 29 mv. negative shift of the half-wave potential with a tenfold increase of the copper-Enta concentration. This was experimentally verified only at the highest buffer concentration compatible with the given ionic strength (see Fig. 1). The half-wave potentials of copper-Enta in solutions where the total ionic strength is that from the buffer components, are plotted against p(H) (see ref. 8) in Fig. 7. A very 500
q
.? 400
t
+
pendence of the half-wave potential upon the buffer concentration.11 The strong shift of the half wave potential toward more positive values with decreasing ionic strength may only in part be accounted for by considering the change of the activity factor in equation 3 and of the K y H n 1 2 values in equation 4, B. Wave Splitting in Unbuffered Slightly Acid Solutions.-In unbuffered solutions of the copperEnta complex the p(H) a t the interface will obviously increase with increasing current, due to the increasing consumption of the hydrogen ions by reaction 2. The relation between the interfacial hydrogen ion concentration (H+),, and the current can be expressed in terms of the concentrations and diffusion coefficients of all proton donors, water included, assuming the validity of the Ilkovic equation, according to d
+
K[DH'/z((H+)- (H+)o] DcUy~'/z((CuYH-)( C U Y H - ) ~ ~ DoH'/~((OH-)O - (OH-)Jl (6)
+
This equation results from the condition of stationarity which presumes that under equilibrium conditions the diffusion fluxes of all proton donors have to equal those of the proton acceptors,13 The first and second right hand members of equation 6 represent, respectively, the Ilkovic fluxes of free hydrogen ione, and of the copper-Enta acid (ses ref. 4); the third term represents the contribution of water (as proton donor) in terms of the hydroxyl ions flux (see ref. 13b); r2 is the mean number of hydrogen ions bound per ion of Enta; it varies with p(H) (see ref. 8) according to Fig. 8 which was computed using the dissociation constants of Enta acid as determined by Schwarzenbach.8 The laat term in equation 6 can be dropped as long as the interface remains acid. Then the current can be expressed more simply by
7 1 1 / / I/------
$
Vol. 58
1
1300
3y 200
4
5
6
7
8
9 1 0 1 1 1 2
PCH).
Fig. 7.-Dependence of the half-wave potential of the copper-Enta complex upon the p ( H ) in buffered solutions. Ionic strength (0.10) from the buffer components, 1 = 20"; copper-Enta concentration: 2( O), 10-4 M ;I(-@-),10-8 M . Full drawn curves were calculated according to equation 5 using Et = -428 mv. (All solutions contained 5 X lo-&'% methyl orange.)
good agreement with the dependence predicted by equation 5 is found in the p(H) range of 4 to about 10 using the value El = -428 mv. The increase of the half-wave potential in alkaline medium i s not predicted by equation 5 . It might be accounted for by the formation of hydroxo complexes of the copper ion and/or of the copper-Enta complex in alkaline solutions. Equation 5 does not account for the strong de(8) G. Schwarsenbach and €€.Ackermann, H d u . Chim. Acto, SO, 1768 (1848) : value8 reported by these authors are concentration mostants; in the following we shall use the value of YH+ = 0.81 for hydrogen ions a t 0.1 ionic strength. The symbol p(H) will be used whenever the concentration (and not activity) of the hydrogen ions is considered. (9) J. Tornea, CoElecfion Czech. Chem. Communa., 9, 12, 81, 190
(1837). (10) T h e interaction of the copper-Enta complex with hydrogen ions can be neglected an numerically inaianificmt fit pT3 hipher than b ( w e ref, 4 and a ) ,
where K C ~ YisHthe stability constant of the copperEnta acid (see ref. 4) and the diffusion coefficients of copper-Enta anion and copper-Enta acid are assumed to be equal. (11) As pointed out by 0. H. MUller in A. Weissberger, "Phyaical Methods of Organic Chemistry," Part 11, 2nd E d . , 1949,p. 1833,adequate buffering is provided when the concentration of the active partner of the buffer is in a hundred-fold excess over thab of the reducible speoies; the hydrogen ion concentration a t the electrode surfaoe ahould then be within about 1% t h a t of the bulk of the solution. This evaluation is based of courae upon the hypobhesis of instantaneous attainment of all proton equilibria. I n our case a solution of 10-4 M copperE n t a in 0.02 114 1:1 aoetate buffer exhibits a half-wave potential which oorresponds to a concentration of the free E n t a ion (Yc-)o at a p(H). by about two p(H) units r n v e alkaline than in the bulk ef the solution (compare Figs. 1 and 7). (12) KYH- values at 0.5 ionia strength were reported by 8. F. Carini and A. E. Martell, J. Am. Chew. Sac., 76, 6745 (1952). (13) Similar relations are t o be found in the literature. See, Le., (a) E. F. Orlernann and X. M. Kolthoff, ibid., 64, 1044 (1942): (b) P. Ruetsohi and S . Trampler, Helv. Chim. Acta, 85, 1021, 1486, 1887 (1 962).
April, 1954
POLAROGRAPHIC REDUCTION OF COPPER ETHYLENEDIAMINE TETRAACETATE343
Equations 5 and 7 can be solved graphically for (H+)oand allow calculating the current-potential
curve for a particular solution. This curve fits reasonably well the experimental results (see Figs. 2 and 3), considering the approximations introduced by using the Ilkovic e q u a t i ~ n ' ~ , ' ~ and the assumption of instantaneous equilibria. The appearance of the double pre-wave can be explained qualitatively as follows. Let us conmolar solution of copper-Enta comsider a plex at a p(H) of about 3.5 (see Fig. 2). At this p(H) the ;ii value will be about two (see Fig. 8). 2 3 4 5 6 7 8 910111213 Now as the current increases (H+)c,as well as P(W. (CUYH-)~will decrease. At a p(H) = 5.5 both Fig,&-The variation of the mean number of hydrogen will be already negligible as compared with the ions bound per ion of Enta with p ( H); calculated according bulk hydrogen ion and copper-Enta acid concen- to Schwarzenbach*ovaluesof the dissociation constants of trations. Until that point the 'ii value will suffer Enta acid; 1 = 20 , p = 0.1. no marked change and consequently. the polaro- stability constants higher than 1010 (see ref. 4). graphic current will reach a diffuslon plateau. I n nearly neutral solutions, the reaction of Enta As the hydrogen ion concentration at the electrode ions (liberated a t the drop surface) with hydrogen surface will decrease further, Ti will start to dimin- ions can thus be neglected until the concentration ish also, until reaching a new steady value of of the metal ions in the interface drops below a about one in the p(H) range 7.5 to 9.5. Thus the very small value. Under those conditions (10) current should start to increase again after hav- reduces to ing reached the first plateau and attain a second i e K D M , ' / z ( ( M ~-) (Me)o) (11) one a t twice the value of the first diffusion current. This second plateau, however, should be According to this equation, the current should rather indistinct as a t p(H) 7 the current will soon reach a steady value, when the metal ion concenfurther increase due to the OH- term in equation 6. tration a t the interface becomes small in compariThe numerical value of the total height of the son with the concentration in the bulk of the solution, A further increase of the current should prewave should be given by then occur due to the increasing contribution of the second term in equation 10. The usual diffusion plateau of the copper-Enta complex should This equation is in a good quantitative agreement be reached finally. The first diffusion current with the experiment (see Figs. 2 and 3). The should be proportional to the metal concentration, first part of the prewave is found to be somewhat in agreement (see ref. 6) with the experiment (see higher than was to be expected. fig. 5 ) . The shift of the half-wave potential of the preThe wave splitting of course will occur only if wave toward more positive values with decreasing the concentration of the metal ion is smaller than ionic strength (see Experimental results) can be the concentration of the copper-Enta complex. accounted for only qualitatively on the basis of As a matter of fact at higher metal concentration the foregoing discussion. a single wave appears, whose half-wave potential C. Wave Splitting in Neutral Unbutrered Solu- shifts toward more positive values with increasing tions of Copper-Enta Complex by Metal Ions.-If a metal ion concentration. nearly neutral solution of the copper-Enta complex In order to express the potential of the dropping contains metal ions Mecz,"+ capable of forming mercury electrode as a function of the current, stable complexes with Enta, then the following equation 4 should be substituted by reaction should be considered in addition to reac4 tion 2 ( Y * ) o= (MeY)o (YH& = J(H+)o f Y4-
+ Me(2)n*
+
YMe(2)('-fi'- (KMeY)
0
(9)
I n the following we shall neglect the formation of ion pairs between the metal Enta complexes and metal ions, as well as the presence of the Me(z,Y complex in the bulk of the solution which is appreciable only for lead under our experimental conditions (see ref. 2 and 6). Then the current observed should vary according to
K~ey(Me)o) (Y4-)o (12)
Combining equations 3 and 12 one gets the following simple expression for the half-wave potential of the metal induced wave id1 El/,= E1 0.029 log K x e y + 0,029 log 2id v (13)
+
-
adi
where i d l is the diffusion current of the metal induced wave. It follows from (13) that the greater i = KDMMeY1/a((MeY)o - (MeY)} KDy1/a(Y4-)oJ(H+)o the stability constant of the metal ion complex KDMe'/t((Me) - (Me),} - I < D u ' / ~ ( Y ~ -J(H+)o )o with Enta, the better the separation of the metal (10) induced wave. Actually the shift predicted by In all cases studied here the metal complexes have equation 13 is only in a rough qualitative agreement with the experiment. Figure 4 shows a half(14) (a) J. Kouteaky and R.Brdiaka, Collection Czech. Chem. C o w wave potential separation of about: 250, 190, muns., 1'2, 337 (1947); (b) P.Delahay and aoll., J . Am. Chem. Boc., 160, 1.10 and 90 mv. for lead, Innt,hannm, sadmiurn, 1948-1863i ( c ) P. Delahay, ibid., 74, 3497 (1962); (d) w e alao ref. 3.
-
344
HAROLD EDELHOCH AND HUGH S. TAYLOR
zinc and calcium, respectively, whereas the predicted values are (see ref. 4) : 530, 450, 480, 470 and 310 mv. According to equation 13 the half-wave potential of the metal wave should shift at a constant metal concentration to more positive values with increasing copper-Enta concentrations. At a constant copper-Enta concentration the half-wave potential should shift toward more negative values with increasing metal concentrations. Both predictions are fullfilled a t least in a qualitative way (see Fig. 5 ) . The dependence of the half-wave potential on the ionic strength (see Experimental part) can only partly be accounted for by considering the influence of the ionic strength upon the equilibria involved. D. Interference of Reaction Rates in the Polarographic Reduction of the Copper-Enta Complex.It seems that a t least some of the discrepancies observed between theory and experiment can be attributed to the influence of non-equilibrium conditions in the mercury-solution interface (see ref. 3 and 14a, b, e). It was found that nickel ion does not produce the wave splitting observed with all other metals studied here. This indicates the possibility that the equilibrium 9 is not very rapid. This possibility is further confirmed by the fact that for all metals studied here the wave splitting is much less
VOl. 58
clean a t the jet electrode than a t the dropping mercury electrode. As pointed out previously’6 a t the jet electrode only very fast reactions can manifest themselves fully. The strong dependence of the half-wave potential upon the buffer concentration indicates that slow proton equilibria in the mercury-solution interface can also interfere kinetically in the observed phenomena. In addition to reactions 2 one should consider here also the dissociation equilibria of the buffer’6 as well as the possibility of a direct reaction between the acid component of the buffer and the Enta ion liberated in the electrode process (see i. e., ref. 16a). Each of these reactions can be slow enough to disturb the overall process. Finally the rather strong dependence of the halfwave potential upon the ionic strength could be interpreted partly a t least in terms of the BjerrumBronsted concept of ionic reactions: the increase of the ionic strength accelerating or retarding the rates according to the charges of the reaction partners.” (15) (a) J. Heyrovsky and J. Forejt, 2. physik. Chem., 198, 77 (1943); (b) Bee also ref. 2. (16) (a) C. R. Castor and J. H. Saylor, J . A m . Chem. Soc., 75, 1427 (1953); (b) 0.H. Milller, ibid., 68, 2434 (1940); (c) see also ref. 140. (17) See, Le., 8. Glasstone, K. J. Laidler and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Co., Inc., New York, N. Y., 1941, Cbapter VIII.
THE ADSORPTION OF GASES ON CALCIUM FLUORIDE1 BY HAROLD EDELHOCH AND HUGH S. TAYLOR Frick Chemical Laboratory, Princeton University, Princeton, N. J . Received November 937, 1068
Isotherms of argon and oxygen, nitrogen and carbon monoxide, xenon and propane, in the neighborhood of liquid oxygen and nit’rogen boiling points, on sintered and ground calcium fluoride powders have been determined. Marked inflections in the isotherms for argon and oxygen were found. Nitrogen and carbon monoxide conformed t o the BET equation (Type 11). With xenon and propane the data fall on a single smooth curve of adaorbed volume versus relative pressure. The data have been examined in terms of two-dimensional changes of state, horizontal interactions and derived thermodynamic data.
Recently, Young and Tompkins2g8examined and augmented the observations and calculations of Orr4 pertaining to the interactions of gases with surfaces possessing periodic structures, such as ionic crystals. I n a re-evaluation of Orr’s work with non-polar gases on CsI, Tompkins and Young* suggest that adsorption was taking place on the (110) plane of CsI. As adsorption progressed to the more uniform parts of the surface, horizontal interaction effects became significant and the localized (square) array of molecules condensed to an hexagonally close-packed layer. This rearrangement in molecular orientation was associated with a phase change. These events con-
tribute an additional heat term based on lateral interactions to the predominant one occurring between the adsorbed molecules and the adsorbent. If the latter heat effects are sensibly constant, a maximum in the heat curve will be evident. Crawford and Tompkins6 have measured the adsorption of polyatomic molecules on BaF2 and CaF2. They have concluded that the adsorbed molecules formed a localized layer on the crystal planes with their distribution fixed by the lattice structure of the adsorbent. The nature of the adsorbed phase has received considerable attention, notably by Hill,s Gregg,’ and Jura and Harkins.8 The latter have postu-
(1) Abstracted from a thesis by Harold Edelhoch, presented in partial fulfillment of the requirements for the Ph.D. degree in Princeton University, 1947. (2) D. M. Young, Trans. Faraday Soc.. 47, 1228 (1951); 40, 548
(5) V. A. Crawford and F. C. Tompkins. ibid.. 44, 698 (1948): 48, 504 (1950). (6) T.L. Hill, J . Chem. Phys., 14,441 (1946); 15, 767 (1947); 17, 520 (1949). For a critical review of physical adsorption theory see chapter by T. L. Hill, in “Advances in Catalysis,” Academic Preas, Inc., New York, N. Y.,1952. (7) 8. J. Gregg, in “Surface Chemistry,” Butterworths Scientific Publications, London, 1949. (8) 0.Jura and W. D. Harkins, J . A m . Chem. SOC.,68, 1941 (1946).
(1952).
(3) F. C. Tompkins and D. M. Young, {bid., 47,77 (1951). Roy. SOC.(London), (4) W.J. C. Orr, ibid.. 86, 1247 (1939); PTOC. A173,349 (1939); J. K.Roberts and W. J. C. Orr. Trans. Faraday Soc., 84,1346 (leas).
