THB MECE&SISM OF THE PEPTIZATIOB AHB
PRECIPITATION OF HEGATIVB SILVER IODIDE SOLS
fcy
William Joseph Coppoc % B. Sc*
A Thesis
Presented to the Faculty of the
lice Institute of Houston, Texas
in Partial Fulfillment of
the Requirements for the
Degree * of
Master of Arts
The Rico Institute
IPS? CONTESTS
ISTRGBUCTXGH 1
EXPERIMENTAL PROCEDURE
(1) THE PEPTIZATION Of SILVER IODIDE 4
(B) THE ADSORPTION OP HIBRQXYL ION 7
(3) 1HI DETERMINATION OF PRECIPITATION VALDES 10
THEORETICAL DISCUSSION 17
S®IRX 24
BIBLIOGRAPHY 25 THE MECLIAHI&Vi OF THE PEPTIZATION AKB
PRECIPITATION OF MBATIV& SlVm IODIDE SOLS
Recent developments of theoretical colloid chemistry bring to the fore the question of the mechanism of peptization and precip¬ itation of colloidal particles; that is, the character of the double layer, the specific action of some ions in forming this layer, and tho change in the double layer accompanying peptization and precip¬ itation.
This problem has been attacked from several theoretical and experimental viewpoints. Freundlioh (2), working with arsenious sulfide sols, proposed the adsorption theory for the precipitation of sols and also believed at that time that equivalent amounts of ions of different valence were necessary for the coagulation, leiser and co-workers {18, 21, 25} using arsenious sulfide, hydrous ferric oxide, and hydrous alumina sols, found that the adsorption of ions of varying valence need not be in equivalent amounts to effect precipi¬ tation and Freundlich {4} recognized in 1929 that the amounts of adsorption were not equivalent. Weiser {19) la 1951 contrasted the adsorption theory of the coagulation with the solubility theory of
Duclaux (1) and Pauli ill, 18, 13) and proposed a quite complete mechanism for the coagulating action of electrolytes on sols of the hydrous oxides and, by inference, other colloidal solutions.
Kruyt and van der Willigen (8) proposed their « isomorphism” theory in 1928 from the results of their experiments on silver iodide sols. According to this theory, a precipitate will be peptized by, and only by, an ion which will form a salt, with the oppositely 2
charged ion of the crystal lattice, which is isomorphous with the pre¬ cipitate. From their discussion, the term "isomorphous” is taken to mean "will form mixed crystals with”. Sound basis for this theory was found in the work of li. Mare (9) whose results showed that adsorp¬ tion is dependent on crystal structure in such a way that “adsorption takes place when adsorbent and adsorbed substance are isomorphous or of similar crystal habit** Thus, potassium chloride will peptize silver iodide forming a negative sol because silver chloride will form mixed crystals with silver iodide, inter work in this same laboratory
(16, 17, 6, 7) has given results which, in the opinion of iCruyt and co-workers, support this theory, and the results of Versey and Iruyt
(17) apparently contradict some parts of the adsorption theory. A true mechanism has, however, not bean proposed; that is, the state¬ ment is made that ieomorphous ions build the double layer but a picture of how this is done is not given. This theory then, as we shall see, reduces to a special case of the adsorption theory as
Koltboff (5) has classified it in his comprehensive paper on "Adsorp¬ tion on Ionic lattices"•
The apparent contradictions which Verwey and Kruyt observed were: (a) Some ions may be adsorbed on the particle but will not in themselves peptize the particle; (b) precipitation my occur before the maximum adsorption is attained; and (c) maximum adsorption is
^Sloat and lieazies (15) in a more recent study of the adsorp¬ tion from aqueous solution of six salts on PbS, found that the adsorp¬ tion was related much less closely, if at all, to the relative lattice dimensions of salt and adsorbent than to the aqueous solubility of the solute. 3
oometimea obtained long before the precipitation value of the sol is
reached. As examples of these contradictions they cite: (a) hydroxyl
iois% as sell as oxalate and phosphate ions, are adsorbed on silver
iodide but are not in themselves "potential-deterraini ng " ions and ere
therefore not able to peptize the silver iodide; (b) cerium ion,
adsorbed on a silver iodide sol in the presence of a large amount of
hydrogen ion, precipitates the sol long before sufficient cation has
been added to effect the precipitation of the sol; and (e) barium ion
on a negative silver iodide sol attained maximum adsorption at approx¬
imately. 0.2 to 0.4 lailliequivalents of barium per kilogram of sol and
yet the sol was not precipitated until the concentration of barium was
approximately 1.0 to 1.2 milliequivalents per kilogram of sol,
feiser and Milligan (22) have explained case (b) from the
fact that, to the undlalyzod sol with which Verwey and Kruyfc were work¬
ing, the large concentration of hydrogen ions would reduce the zeta
potential to the coagulation value before the maximum adsorption of
cerium ion had been attained. In this study, it is proposed to explain
cases (a) and (e) to the light of the adsorption theory and to propose a mechanism, according to the adsorption theory, for the peptization and precipitation of negatively charged silver iodide sols.
It Is important to observe that in all of these cases, the precipitation value plays quite an important role la determining the validity of any one theory. For example, Fennycuick and Best (14) and
Weiner and Gray (20) have also used the precipitation value as a means of solving problems of theory.
In the experimental portion of this paper there ia given: 4
(1.) The effect of various loss on the peptization of silver
iodide.
(2.) The adsorption of hydroxyl ion on silver iodide*
(3) The determination of the precipitation value of several electrolytes for silver iodide sols.
BgBBIHSna&
(1) FEPTIZATIOH OF SILVER IODIDE.
Since, as Kruyt and Oysouw 16} have pointed out, and as has been found experimentally daring this study, only the iodide ion will serve to peptize directly a well-washed silver iodide precipitate, it was decided to add the peptizing electrolyte to the potassium iodide solution before mixing the two solutions to form the sol* By this method, the final sol always contains a known concentration of potassium nitrate. The slight solubility of silver iodide is suffi¬ cient to allow quite rapid aging so that by the time the silver iodide, which has been precipitated from equivalent solutions of potassium iodide and silver nitrate, has been washed free from electrolyte, the aging of the particles has proceeded to a point where direct peptiz¬ ation is almost impossible*
The solutions of silver nitrate and of potassium iodide were prepared from the "C. P.* salts and were found to contain no impurities which mi^ht interfere with the accuracy of our results. The solutions were made up approximately 0.05 Jf and standardized, then carefully diluted to 0,0223 H.
Approximately 1.0 I? solutions of electrolyte were made up by 5
placing a known amount of the properly dried salt in water and dilut¬
ing to 100 cc. except in the case of potassium hydroxide which ©as made
up approximately correct and then standardized against standard hydro¬
chloric acid. The proper amounts of these solutions to give the
different concentrations of electrolyte desired ©ere added to enough water to make 10 cc. and this solution then added to the potassium
iodide solutionj thus in each case, 48 cc. of silver nitrate solution,
45 cc. Of potassium Iodide solution plus 10 cc. of the electrolyte solution were mixed. Since the silver nitrate and the potassium iodide were both 0.0S22 H, 100 cc. of a sol containing 10 mMola of Agl per liter was obtained.
The solutions were mixed In a "quick-mixing" device of 100 cc. capacity. This device consisted of a 50 cc. Pyrex beakex* placed inside a wide mouth glass chemical container with a bakelite, screw- top lid and held in place by a glass stirring rod which was of the proper length to reach exactly from the bottom of the beaker to the opposite side of the screw-top lid. The silver nitrate solution was placed in the beaker and the potassium iodide solution plus the peptizing electrolyte was placed in the container outside the beaker, the lid screwed on tightly and the whole then shaken vigorously for thirty seconds. This device assured probably the quickest and most complete possible mixing and was found to give easily reproducible results after a little practice.
The electrolytes used were potassium bromide, potassium fluoride, potassium nitrate, potassium chloride, potassium sulfate, and potassium chlorate. According to the isomorphism theory of Kruyt, Peptization Properties 2 0 o S3 O o o o o SO ♦ * * *r4 H •*4 «H *0 *d 4* 0 % *0 #2 4» 49 £) % •P m 0k A & 0 n* 0 04 m 0 Pi 23 M 0 49 04 fil % m a 0 A 5> a, 0 * ts3 m “ a _ p i-4 *0 49 49 4> h m 0i 0 0 0 «$ «*4 «H 49 0 <& 4> 49 *H «rt nd -nd ^4 jSt *rl i~* 4> ,0 P 0 A P* Pt Pi Pi P§ A 0 m Pi 0 rk © © 4* o 0 v4 fa t* A |ft N§ 0 N 0 0 I P* 0 & 0 PU 0 M s* 49 4» JH % 9 0 0 0 *H i4 *g ** 49 0 *r4 «H *d *0 *0 4^ 4> 49 m 49 JP ns ** 4» tJ H A 0 0 v« 0 aa iS2 A Oi 0 5* Pi 0 m 0 4* 04 A Pi A 0 0* > 0 * 0 2a *a 0 «D $2 0 . 0 m 4> % 0 p 0 *0 v4 *4 i4 **4 4» *d *a 4» 4> *0 »H i*4 42 4» 42 *H *Hf •*4 %4 4* *H *rl oat *d 49 4* 4» 4* 0 S* H a 0 th Pi 0 A a 0 <8 0 0 m 0 m 0 0 O Pi a. 0 0 0 0 O o m & % A P4 0 0 0 *4 H m 5s *Cf <0 -H *ri T4 *rt 42 *0 m *d *ts 42 H «*-t *r4 f4 n-4 «r4 49 42 un i*4 if4 *r4 i4 4» 4>- *xi *0 4» 42 49 U fc 0 Pi 04 A 0 O 0 0 0 0 0. 0 0 . 9 Pi A m P* 0* 0 m O 0 A 0, a 04 D< 0 m 0 0 Q m m o 0 M £$ •Ht «r4 •HI H i4 *4 *0 42 49 *4 *#4 4* 42 42 n*$ «H *4 i4 49 $$ *H *4 *0 m *rl i4 «ti ts P 49 m h U 0 h 0 Pk A 0 O 04 04 Pi 0 m 0 aa 0 % h © 0 a o m 0 Pt 0. 0 0 m o 0 m 0 fj m P* m aa P* 3 m 8 M 29 Iianefiiately precipitated precipitated precipitated precipitated After IB Hoars precipitated precipitated precipitated precipitated 6 7
only potassium iodide, potassium bromide, and potassium chloride
should be able to peptize silver iodide. This was found to be true.
In the first series of determinations, all of the above salts
except potassium hydroxide were used, each electrolyte being tested at
concentrations of electrolyte equal to 0.1, 0.05, 0.G25, 0.005 H. The
results of this series are summarised in Table I. The nature of the precipitate in those case© where the precipitate had settled out over night was decidedly different from those in which coagulation occurred
immediately on mixing. In the former, the precipitate was of a fine
powdery nature and in the latter it was in the form of the familiar large floes.
U) THE ABSORPTION OP BYDftOm ION.
It was next proposed to make a series of determinations using potassium hydroxide as the peptizing electrolyte. The same con¬ centrations were used as in the former trials. At the end of twelve hours, only the material in the 0.025 N solutions was found to be peptized although the precipitate in the other solutions was largely of the powdery nature found in the 0.1 H solutions of the halides. It was decided to investigate this interesting case more thoroughly both to determine the limits of this peptization range and to determine the exact fora of the adsorption curve.
The peptization range was determined by decreasing the con¬ centration difference between samples and the total concentration range of successive runs. In each case where precipitation had occurred, the supernatant liquid was analyzed for free hydroxide ion by titrating the supernatant liquid with hydrochloric acid. The results with regard to Mo/s. L/fer 8
Table II
Peptization Range of SOB done. Time Iquiv. after per mixing Trial 1 Trial 2 Trial 3 liter Hours
0.000 00 precipitated precipitated precipitated 12 precipitated precipitated precipitated 0.003 00 precipitated 12 precipitated 0.010 00 partly peptized 12 precipitated 0.080 00 nearly all peptized peptized 18 slightly peptized precipitated 0*081 00 peptized IS precipitated 0.028 00 peptized 12 precipitated 0.023 00 peptized 12 partly peptized 0.024 00 peptized 12 peptized 0.025 00 peptized peptized 12 peptized peptized 0.026 00 peptized 12 peptized 0.027 00 peptized 12 peptized 0.028 00 peptized 12 slightly peptized 0.029 00 peptized 12 precipitated 0.020 00 nearly all peptized peptized 12 slightly peptized precipitated 0.040 00 slightly peptized 12 precipitated 0.050 00 precipitated precipitated 12 precipitated precipitated 0.100 GO precipitated 12 precipitated TSabie III
Adsorption of %droxyl Ion by Agl
Original Equilibrlna Adsorption Cone* in Oono. In Or. Qir per Equiv./lit. Iqnif./lit, Or. Agl
O.COSO 0.00408 0.0070
0.0200 0.0188 0.0087
0.0260 0.0S49 0.0080
0.0270 0.0261 0.006S
0.0280 0.0273 0.0050
0.0260 0.0284 0.0043
0.0300 0.0290 0.0078
0.0400 0.0388 0.008?
0.0500 0.0485 0.0110
0.1000 0.0981 0.0137 10
peptization and adsorption are shown in Tables II and III, respectively*
The shape of the adsorption curve is shown in Fig. 1 and sill be dis¬
cussed somewhat in the theoretical portion of this paper* Each point of the curve was determined at least twice. It was found that the exact limits of the peptization range were very sensitive to slight excesses of either silver or iodide ion in the original solution; being widened somewhat by an excess of iodide ion and giving a slight brown coloration of silver oxide if a slight excess of silver ion were present in the solution*
The supernatant liquid from those solutions above the peptiz¬ ation range were analyzed, qualitatively, for free iodide ion. We neutralized these solutions with nitric acid end added silver nitrate but were unable to detect any iodide ion in any of these solutions.
(3) DETiesmTICfl OF PfiEOIFIThTTOH mOSS.
An attempt to measure the adsorption of cations by measuring the amount of hydrogen ion displaced by the addition of electrolyte to an electrodialyzed eol which had been concentrated by evaporation ms not found to be feasible. The change in pH between the original sol and the coagulated sol was too small to be significant. Verwey and
Kruyt (loo, olt.) also found this to be true and attempted to circum¬ vent the difficulty by making their measurements on the ultrafiltrate.
This procedure was not used since it was thought that the increased change in the pH was in itself evidence that the adsorption equilibrium had been displaced. It was during these preliminary experiments that the large bulk of precipitate which formed in those solutions which were considered to be just below the precipitation value was noticed 11
and it seemed that the difficulty might be attributed to this fact.
Freundlich (3) has defined the precipitation value as "that one con¬ centration of electrolyte which will develop a given change in a sol within a given time, i.e., the blue color in a red gpld sol within five minutes, or the complete precipitation within one or two hours of an iron oxide or arsenic trisulfide sol**, fhis definition is, of course, perfectly satisfactory when one is comparing the action of different electrolytes on the same sol, but when one is comparing the precipitation value of the sol with some other property of that sol, then one must consider more carefully just what will be called the precipitation value. In this case, fear example, the supernatant liquid might be still nearly as opaque as that in the original sol and yet the large amount of eoagulum was evidence that most of the sol had precipitated, for the double reason then of seeking an explanation for the abnormally high preoipitatlon value observed by ?erwey and Kruyt, and of studying the precipitation value for its own sake, it was decided to investigate the phenomenon by the procedure which follows.
THE PREPARATION, OF THE SOLS: A very concentrated sol was thought desirable for this work in order to make the differences as large as possible. Solutions of silver nitrate and hydrogen iodide were made up approximately G.2G j| and G.22 H, respectively, titrated against each other by neutralizing the hydrogen iodide with excess calcium carbonate and taking the end-point with potassium chromate as indicator, and the proper amounts of the two solutions mixed in the quick-mixer to give about 500 co. of a sol containing approximately
100 mMole of Agl per liter with 10$ excess iodide ion as the stabilizing 12
ioxu The quick-mixer was of the type previously described except that
it had a ground glass lid and was large enough to mix 800 cc. at one
time* The solutions were put into the mixer and the whole then cooled
In an ice bath for half an hour or longer, the mixer was then removed and shaken for thirty seconds to insure complete mixing and the sol
then iiss&edlately put in an electrodialysia bag and electrodialyzed for
48 hours at 700 volt® and 16-20 milliamperes* The dialysis bags were of cellophane which had been previously soaked in 63$ ZnClg (10) and
then washed with distilled water to remove the electrolyte. After the
48 hours dialysis, the relatively pure sol was evaporated to 100 cc* by passing a current of cleaned and dried air over it* At the end of
this treatment, the eels were quite stable, had a pH of 3*5 to 3*7, and contained approximately 60 grams of silver iodide per liter. When
1200 ce. of sol bad been prepared in this way, the precipitation experi¬ ments were commenced.
In the precipitation experiments, 15 cc. of sol and 5 cc. of electrolyte solution of the concentration desired, were mixed in a small, all-glass '’quick-mixer” and shaken for thirty seconds, the solution then being transferred to a Pyrex test tube, allowed to stand for IB hours, and shaken again for ten seconds. The purpose of this 18 hour shaking was to break up the sort of gel-like structure which must have existed
in the solution since this shaking always resulted in more of the material being thrown down. Subsequent shakings seemed to lave no effect. At the end of 24 hours, a 5 cc. sample was pipetted from the top of the solution and analyzed for silver iodide by completing the precipitation with barium chloride, filtering through a Gooch crucible, O SoJ. 1 Cloudy Clear
Normality x 10z Norm o/z'-f-y x J O So/. 1 Cloudy C/ea
Normality xjot Norm al/Jy x/O5 13
and sashing with very dilute nitric acid (about 0.5 cc. of cone, reagent to 200 ec. of sash solution) until the washings gave no pre¬ cipitate with sulfuric acid. The crucible was then dried over night at 110°C. and weighed.
A complete table of the run using lanthanum nitrate on sol 1 is given in Table IV. Tables V and VI contain the results for the other electrolytes in a condensed form. These results are illustrated graphically in Figures 2 - 7. In each case except that of barium chloride on arsesious sulfide sol, tea solutions were mixed covering the precipitation range and samples removed from as many as was thought necessary to give a complete picture of the curve. In th© case of the barium chloride on arsenious sulfide, only eight solutions were mixed because of a shortage of that sol. One solution was always made which was diluted with pure water and that solution used as the basis for calculating the percentages of that particular curve. In each case, the amount of electrolyte solution added to enough water to make 5 cc. ranged from 0.0 to 4.5 or 5*0 cc., usually in 0*5 cc. steps* The electrolyte solutions were made up slightly stronger than was thought necessary and then diluted if it was found that they gave precipitation at values too low for convenient measurement. The solutions were analysed by standard methods with the exception of the lanthanum which was prepared by weighing out a definite amount of lailOgJg.SBgO. the potassium chloride was analysed by precipitating the chloride as silver chloride and weighing as such; the calcium sulfate and the calcium chloride by precipitating as the oxalate and igniting, then weighing as the oxide. The aluminum chloride was precipitated and 14
Usable IV
Solution Cone, of &n’t of Percent Condition of Humber Ia(ll03}3 Agl per Precis}* d Supernatant Equiv./Ut* S- ce. of liquid Sup’usft x io4 Liquid . Grass
1 8.10 ««#*## clear
2 0.30 clear
3 4.89 0.0000 100.00 clear
4 4.28 #*»*#* very slightly cloudy
5 3.6? 0.0003 99,86 slightly cloudy
6 3.06 0.1164 47.04 nearly opaque
7 S.4& 0.2073 5.89 opaque
a •' 1.84 0.2164 1.56 opaque
9 1.23 opaque
10 0.00 G.81SS 0.00 opaque The Effect of Various Electrolytes on the Percent Precipitated and the Opacity of the Supernatant Liquid of Agl Sol 1 a •H to lO 02 to o> CO © H H *P O o © © al 02 4 4 rH CD rH •O rH O O H 8 o 02 tO rH o© rH H H 02 'HQ U 00 CD O H OH a JU tO 02 CO 00 H H CO O 02 o- o o o o 02 rH o • © • Pi • • • ■ • ♦ • • « • • • rH CD 02 • o • a> • 3 • H • U : 8 • © * 4 4 4 *P -p H H *a *d Pi © 1 © o o P« O © O O P © P P 2 © i» © Pi >> a o a © p Pi © o • • • • • •H *d & 4 H +> 4 4 O H«C0 to OH £N O© to o woo rH H CD• 02 O 02 OrH o p tN 00 oo to»d H CD• H CD• H i—1O Hio-d rt ^Pi O ©H H CDO 02 IQO ts H«d CD OH o 00 HO o o • Pi-H • © • © • • H • H ♦ •P • ♦H • • • • ♦ o O 00 H o u rH O cd 9 >• to >» O >> 0* K*) 02 H • O ♦ rH ♦ *d • • 4 8 © }» © © © © O o p Pi © o • • • • g. 4 4 Bl 4 H -P H O 00 OH O © rH •♦ H •O CD rH 00 rH CD IS CD CD H IS H CO H O 0202 00 OH CO 02 to • o 1> H CD ♦ o o o ss o © P'sR Pi O ©«H • «© ♦ • • • • • • • • ♦ • rH O O H O U o CD to 02 to © Pi UH „ o' • rH • H • • * • • 4 4 4 -P rH nd *d np © «d ri © © O o >> o O O o © p © P o P Pi © Pi © p © >> © Pi © a a a • • • • • ♦ al 4 4 4 O 02(52 CD CO CO 3 00 a§ 02 to 02
^Slightly cloudy ♦I,'early opaque 16 17
ignited, then weighed as the oxide.
THEORETICAL BISCUSSIGK
In a paper which has come to our attention since completing the experiments on potassium hydroxide, Kruyt and Cysouw (6) discuss a phenomenon similar to the peptization of silver iodide by potassium hydroxide with regard to the peptization of silver iodide by oxalate and by phosphate ions. These experiments will be discussed in con¬ nection with our discussion of the peptization by potassium hydroxide.
It is evident fro® Fig. 1 that the adsorption of hydroxyl ion follows a perfectly normal adsorption isotherm up to the peptiz¬ ation range, seems to drop somewhat through the range as the particles coalesce and the surface decreases, then resumes a normal course on the upper side of the peptization range. The curve through the range is naturally rather indeterminate both because of the difficulty of obtain¬ ing true measurements in the range and because of the extreme narrowness of the range. However, it is to be expected that the adsorption imme¬ diately above the range would be slightly smaller than the adsorption just below the range because of the coalescence of the particles and the corresponding decrease in the total surface. Mo satisfactory con¬ clusions as to the significance of the adsorption curve with respect to the peptization range have yet been reached, although further study nay show that the theory advanced by feieer (18) to explain a similar case say also function here. That is, that the adsorption of the cation is comparable to that of the anion and at this range only do the curves become far enough apart to permit peptization. 18
IDruyt and Verwey (1?) have this to say concerning the peptiz¬
ation power of potassium hydroxide for the silver iodide sol: "One can
cose to false conclusions if on© adds the peptizing ion to the potassium
ion in the first place; through co-precipitation of otter silver salts,
one could set free a corresponding amount of iodide ion which will then peptize the silver iodide in any case. Addition of the potassium hydroxide to the potassium iodide solution end subsequent mixing with an equivalent amount of silver nitrate leads to complete peptization;
it can however be easily Shown that the sol is only apparently stable,
...Although the solubility product of silver hydroxide (10 ) is not reached in the final condition, it sill naturally be present in con¬ siderable amounts, it precipitates itself little by little around the silver iodide... Hydroxide ion therefore will not peptize the sol, still it will be strongly adsorbed”. They resort to the same type of device to explain the peptization of sliver iodide by phosphate and oxalate ions; that is, some silver phosphate or silver oxalate is formed, thus releasing iodide ion to the solution to peptize the pre¬ cipitate. Their proof for this theory consisted of the fact that they found phosphate ion in the coagulura and iodide ion in the ultrafiltrate after ultrafiltration of a sol which had been peptized by addition of phosphate ion to the solution. Although we did not find any iodide ion in the supernatant liquid of solutions even much higher than the peptiz¬ ation range, we would hesitate to state that there was none there since we did not concentrate the solutions before testing for iodide and since it is always possible that an analytical method is not sufficiently sensitive to detect a very small amount of a substance which might still So/ 1
of' N or m a/ifies. IS
be present.
Verwey and Kruyt mention that the solubility product of . silver hydroxide is not attained in the concentrations of hydroxyl ion studied. A fen calculations will show that if one is to displace sufficient iodide ion to peptize the sol (a concentration of 10 is necessary according to Kruyt and van der Willigen (8)), and if one has a concentration 0*1 H with respect to hydroxyl ion, then one will still have leas than the solubility product of silver hydroxide by a factor of 104. It is of course still possible that in the method of mixing used, some silver hydroxide was obtained as a co-precipitate with the silver iodide and their hypothesis might still be correct* However, if some of the hydroxyl ion is used to form silver hydroxide, it seems that the observed "adecrption" of hydroxyl ion has already been ex¬ plained. The results obtained seem to show that the hydroxyl ion is truly adsorbed and does truly peptize the Bilver iodide sol over the rather narrow range shown.
In the experiments of Kruyt and Cyeouw (6) using phosphate and oxalate ion one has a similar ease. Their experiments showed the presence of phosphate in the coagulum. This might mean either the formation of aggPC^ as Kruyt supposes, or it might mean merely the strong adsorption of the phosphate ion in the true sense. It is very difficult to determine from qualitative tests of tails nature whether the results era due to adsorption or to compound formation. It aright be possible to clear this matter up by studying the adsorption curve of phosphate ion on Agl sols. If the curve is a straight line, then compound formation occurs} if the curve is of the well-known adsorption 20
Isotherm type, then it is an adsorption phenomenon. It will be noticed that the adsorption studies with hydroxyl ion gave a typical adsorption
isotherm. It is to he regretted that a study of the adsorption of oxalate and of phosphate on silver iodide has not been attempted*
In all of the above, phenomena have been dealt with which occur at the surface of the particle. It is well known that the sur¬ face of the silver iodide particle consists of ’•active’* pointe which are points at which the lattice is supposed to he incomplete. If one should prepare a sol in excess iodide ion so that the sol is peptized by adsorbed iodide ion forming the inner layer of the double layer, and then add phosphate ion to this sol, some of the univalent iodide ion would be expected to be displaced from the inner layer by the trivalent phosphate ion. This phosphate ion would remain in part on the particle and would be detected in the eoagulum even after ultrafiltration. It would in this case actually be the peptizing ion and this displacement could hardly be called compound formation. If equivalent quantities of silver nitrate and hydrogen iodide were mixed in the first place, then the small amount of iodide ion which is always present on the surface of silver iodide which has been precipitated from ’’equivalent” solutions
(5), plus the iodide ion which might have bean displaced from ’’active" points on the surface of the particle, would be sufficient to give a test for iodide ion in the ultrafiltrate, especially when this ultra¬ filtrate is subjected to concentration before testing. Thus iodide would be detected in the filtrate and phosphate in the coagulant. It seems that these testa therefore might just as well be used to prove the adsorption theory as to show compound formation* C/ouc/y C/ejzr
Norma/ify x /O3 81
A study of Figures 8, 3, 4, 5, and 7 will aims that on©
may easily be deceived if one considers the precipitation value to
be that value at which all of the sol is precipitated. It is very
evident that for the comparison of the precipitation value of a sol
with some other property of the sol, another definition is necessary.
That is, when one considers the precipitation value in the light of the
old definition, instead of determining the precipitation value for the
large percentage of the mass of a polydiaperce sol, one is in reality
determining that value plus the large mount of electrolyte which is necessary to throw down the very mall particles that constitute only a very small percentage of the mass of the sol. The actual adsorption
by these particles would be very small in comparison to the large amount of adsorption by the large percentage of the mass of the sol, even
though this larger percentage is made up of larger particles with less
surface per unit of mass.
Although Verwey and Kruyt did not define the precipitation value which they used, it is logical to assume that they used the defi¬ nition of the value which has been commonly used up to this time.
This accounts for the abnormally high precipitation value observed by them.
A consideration of just what one should call the precipitation value of a polyd isperse sol leads one to consider the purpose for which it is to be determined. As has been stated previously, if one wishes only to consider the relative precipitating powers of different electro¬ lytes, then there is no reason shy the old method should not be used as long as one experiments on the same sol. If one is comparing the 22
precipitation value with some other property of the sol, then one must consider that it is, in general, the large percentage of the mas of the sol which is determining the magnitude of this other prop¬ erty. For this reason then, one should consider the precipitation value to be that point at which the large percentage of the mss of the sol is precipitated. Several possibilities presented themselves here as means of determining this point from the curves of Figures 2-7. (a) One might extrapolate the nearly vertical part of the curve to the 100$ line and consider this to be the value; however, this value will always be slightly high; lb) one might draw a line through the mid¬ point of the vertical portion of the curve and parallel to the "percent’* axis and consider this quantity to be the precipitation value; this value ie seen to be always slightly low; (c) one might extrapolate the nearly vertical portion of the curve and the nearly horizontal upper portion of the curve and consider the point of intersection of these two lines to be the precipitation value. This last possibility seems to us to be the best quick method since it will always be very nearly correct. All three of the values would bo very close together on the curves obtained in this study and the importance of these curves lies not so much in the determining of a new definition of the precipitation value as in the illustrating of the large "spread" between the point where, say 90$ of the mass of the sol is precipitated, and the point where the solution is clear* Thus it is seen that it takes one and one-half to three times as much electrolyte to precipitate all of the sol as to precipitate ninety percent of the sol; with very much more S3
concentrated sols, the ratio would be even higher.
Finally, it seems that the mechanism of the precipitation of silver iodide hydosols is one which ie entirely conformable to the adsorption theory* Kruyt himself (7) explains the peptization at last on the basis of an adsorption of iodide ion. Although providing a good basis for predicting the adsorbability of an ion when other factors are not too strong, the "isomorphism" theory does not lend itself handily to the formation of a mechanism or complete explanation of the peptization and precipitation process.
Naturally, any theory of peptization must be reversible in the sense that it must at the same time lend itself to the explanation of precipitation. Hence, the mechanism of the peptization must be the building up of the charge on the particle through the adsorption of the negative ions and the consequent increase in the thickness of the double layer until the charge on the particles is sufficiently great to prevent their coalescing. On the other hand, the displacement of the ions forming the outer layer either by ions possessing a higher valenee or by a larger concentration of the counter ion already present will tend to make the double layer thinner, thus decreasing the charge and effecting coagulation.
The question of the cause for the "specificity" of some ions for the building of the double layer is still open although several factors enter in, such as relative crystal structures, solubilities, etc. saasam
I. ¥©rwey and Kruyt found suae ions which were strongly adsorbed by Oliver iodide but which would not peptize silver iodide* These results have been explained from a consideration of the adsorp¬ tion curve of hydroxyl ion on silver iodide and the experimental results and procedure of Verwey and Kruyt by shoeing that, although the conclusions sere in some respects contrary to the adsorption theory, the actual results obtained were not, II. A rather narrow peptization range of hydroxyl Ion for silver iodide has been observed. It is possible that this range say be attributed to coincident adsorption of the counter ion. III. The abnormally high precipitation value, observed by Verwey and Kruyt, of barium ion for silver iodide sols has been ex¬ plained from the fact that the old definition of the precipitation value gives too high a result for polydisperse systems of the type of the usual silver iodide sol. I?, A new method of obtaining the precipitation value for polydisperse sols from a curve of percent precipitated plotted against concentration of electrolyte has been proposed. This new method la usually faster and gives a result which is more nearly true than the old method. ¥. A simple mechanism for the peptization of a negative AgX sol has been proposed on the basis of the adsorption theory* BIBLIOGRAPHY
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