A X A L Y TI CAL EDI T I O S
78
Vol. 2, No. 1
The Problem of Dilution in Colorimetric H-Ion Measurements 11-Use of Isohydric Indicators and Superpure Water for Accurate Measurement of Hydrogen-Ion Concentrations and Salt S. F. Acree3 and Edna H. Fawcettd B U R E ~ OF U STANDARDS A X D BUREACOF PLAXTINDUSTRY, WASHINGTON, D C.
It has been shown that precision pH data in very diHE accurate determiThe isohydric i n d i c a t o r lute or weakly buffered solutions can be obtained nation of hydrogen-ion t e c h n i c gives the correct colorimetrically only by (a)the adjustment of the pH concentrations or activicolorimetric pH of the nnof the standard indicator solutions, (bj the use of the known solution w h e n t he ties in very weakly buffered isohydric indicator methods, (c) the use of superpure latter does not change the solutions (8) is very iniporpH 7.0 water, and ( d ) the correction for the salt and tant in many industrial and color of the added indicator protein errors. scientific problems, such as solution. Tn other words, A method is given for the adjustment of the pH of the refining of sugars, the if the unknown solution and the indicator standards against 0.001 Mor 0.05 M bufferuse of salt in curing green the indicator solution have color standards. hides, the bleaching and purithe same p H value, or same The technic for the use of the isohydric indicator fication of cellulose fibers, hydrogen-ion concentration method is described. or activity, and are suffit h e changes in dilute The production, test, storage, and use of superpure ciently dilute to be free from bacterial culture media, and water in precision pH work are described. salt and protein errors, they the preparation of s t a b l e Tables and curves are given to show the errors in pH colloidal s o l u t i o n s . I n are said to be isohydric caused by dilution of buffers with ordinary pH 5.0 many of these cases comparaand preserve this p H value distilled water and with pH 5.7 air-C02 equilibrium tively neutral materials are when mixed in any ratio. water. In the use of this method washed or dissolved in variElectrometric and colorimetric data are presented to to measure the colorimetous types of n a t u r a l o r show the relation between the dilution of buffers and d i s t i l l e d w a t e r s and the ric pH of the more conthe salt and protein errors of the indicators used. centrated solutions, howp H of such solutions must ever. the color of the indibe determined. It is often desirable to use the colorimetric method with fairly high cator is considered a function both of the electrometric precision in such work, but heretofore unknown factors p H and of a salt error dependent upon the concentrain both the equipment and method caused discordance tion and character of the salt and ions present. The between the p H values determined electrometrically and salt and ions in the more Concentrated solutions doubtless colorimetrically. For example, a concentrated solution of affect the activity of hydrogen ions and electrode material, commercial sodium chloride used for curing hides not only and therefore the e. m. f. reading, but this is used to calcurendered a hydrogen electrode inoperative, but also gave late the electrometric p H taken as the standard pH. The erroneous colorimetric data unless corrections m-ere made salt and ions may affect the color of the indicator through for salt errors amounting to 0.3 to 0.5 pH. Likmise, a re- formation of double salts and otherwise and also through distilled water of high purity, showing approximately the changes in the activities of the hydrogen. hydroxyl, and other theoretically correct p H 7.0 when tested with a p H 7.0 bromo- ions in the solutions involved, and of the molrcules and ions thymol blue solution, gave the two erroneous readings p H of the indicator. Both effects, however, classed together 6.4 and 7.5, respectively, when tested with p H 6.2 and 7.6 as a “salt error,” are added to the factor for the electrometric p H of the concentrated salt solutions in giving data for bromothymol blue solutions. These and many other examples in the accompanying calculating the concentration or activity of the indicator tables show that for accurate colorimetric measurement of ions responsible for the color and the colorimetric pH. The hydrogen-ion concentrations or activities in dilute solutions object of these investigations was to differentiate between or in weakly buffered solutions of high neutral salt content, these factors and measure them precisely. The usual effect attention must be given to three important factors: (1) the of neutral salts on the sulfonphthaleins, and on methyl red adjustment of the p H values of the series of solutions of each above p H 5.4, is equivalent to a decrrase in hydrogen-ion concentration or activity or an increase in colorimetric pH. indicator and use of the isohydric indicator method-that is, defining the colorimetric p H of the unknown as that of A concentrated salt solution producing no change in color the adjusted indicator solution whose color is not changed in the added indicator is therefore colorimetrically isoactive thereby; (2) the prrparation and use of very pure water with the indicator solution, but cannot be considered strictly in making solutions; (3) the correction for salt errors found isohydric therewith. For practical purposes, however, it in the more concentrated buffer-color standards and un- seems best to use the familiar term “isohydric” in this article dealing with dilute solutions for which the terms “isoactive” knowns. and “isohydric” become identical when the salt error beLReceived November 7 , 1929. Presented under the title “Use of comes zero. By definition, the difference between the elecIsohydric Indicator Methods in the Preparation, Tests, and Storage of tromctric p H values for a solution of indicator and a solution Superpure Water” before the Division of Water, Sewage, and Sanitation Chemistry at the 77th Meeting of the American Chemical Society, Columbus, of salt and indicator which are colorimetrically isoactive is Ohio, April 29 to May 3, 1929. the “salt error;” that is, the uncorrected colorimetric p H of a 2 Published by permission of the Director, National Bureau of Standards concentrated solution contains the salt error, whereas the cora Principal scientist (chemist), Bureau of Standards. rected colorimetric p H is identical with the electrometric pH. 4 Junior pathologist, Bureau of Plant Industry.
T
I N D CSTRIAL A N D ENGINEERING CHE;MISI’RY
January 15, 1930
An attempt is being made to analyze the electrometric pH values into the component parts related possibly to (a) the effect of the salt and ions on the electrical activities of the different types of electrodes themselves; (b) a subsidiary positive or negative e. m. f . produced directly by the salt and ions themselves; ( c ) the activity of the hydrogen ions alone, which varies with their concentration and the salt content; and (d) correct measurement and calculation of the liquid junction potentials. rill of these effects may vary when measured by different electrometric, colorimetric, reaction velocity, and other methods. Increase in the concentration of neutral salts generally decreases the electrometric pH but increases the colorimetric pH; the same salt sometimes increases and in other cases decreases the velocities of various catalytic reactions. Adjustment of pH Values of Standard Indicator Solutions By making up a series of 0.001 iZI solutions of each sulfonphthalein indicator such as bromothyniol blue and adjusting them in steps 0.2 pH apart over the useful zones, pH 6.0 to 7.6 in this case, any unknown solution with a pH falling between pH 2.0 and 10.0 can be tested by means of this indicator series. Table I-Adjustment
1
of I n d i c a t o r S o l u t i o n s
COLORIMETRIC
PH OF 0.2
ELECTROMETRIC PH 0.2 ml. 0.001 M bufferb M BUFFER^ compared with 0.2 ml.
OF 0,001
indicator plus 0.2 ml. 0.05 M buffere Bromocresol green 4.5 5.3 Bromothymol blue 6.1
4.3 5.05
6.3 6.5 6.7 6.9 Phenol red 7.3 8.3
ML. INDICATOR PLUS:
0.05 mi. 0.001 M bufferb used in stepwise adjustment and compared with 0.2 ml. indicator plus 0.2 ml. 0.05 ‘24 bufferc
6.05
1 1
6.2 6.4 6.6
6.35 6.45
7.05 8.1
a Buffers made b y dilution of 0.05 M buffers, 1 ml. t o 49 ml. of C o r free water. T h e change caused by this much dilution of t h e series falling between p H 4.2 a n d 8.0 was found electrometrically t o be about 0.25 p H . b 0.001 M buffers having the electrometric p H values given in column 1. c 0 05 M buffers ha7ing t h e same electrometric p H values a s those given in column 1 for t h e 0 001 Ad buffers.
The adjustment of the pH of the standard indicator solutions to eliminate their salt errors was done as shown in Table I, and may be illustrated with bromothymol blue. The indicator was dissolved by warming it with one equivalent of alkali, preferably in a 0.05 N solution, and diluted to 0.04 per cent, or better to 0.001 M, as desired (8). The stock solution was first divided into several 100-ml. portions, and the adjustment to the series of desired pH steps was begun. A volume such as 0.2 ml. of one 100-ml. portion was put into a Pyrex test tube. Another 0.2 ml. v a s put into a second Pyrex test tube to which Jvas added 0.2 ml. of a 0.001 A1 (or 0.2 nil. 0.05 -11)(8)phosphate buffer solution having the electrometric pH 6.2; this small volume formed the first provisional “micro” buffer-color standard. The two test tubes were stoppered and laid on their sides over white porcelain or paper so that the colors might be compared through substantially equal areas. As the hue and brilliance of the sulfonphthaleins vary greatly with their pH, it was readily seen that the stock solution had a pH value below that of the “micro” buffer-color standardnamely, about 5.9. Accordingly, the 100-nil. portion of stock indicator solution was titrated Rith 0.05 -V alkali until a 0.2-ml. portion thereof exactly matched the color of another
79
0.2-ml. portion thereof plus 0.2 ml. of the pH 6.2 0.001 M buffer solution. This 100-ml. portion of the indicator was then set aside as the standard pH 6.2 bromothymol blue solution. I n order to adjust a second 100-ml. portion of the stock indicator to pH 6.4, 0.2 ml. each of the p H 6.2 indicator and a pH 6.4 0.001 iM phosphate buffer solution were mixed. This made a second provisional “niicro” buffercolor standard of approximately pH 6.4 against which 0.2ml. portions of the second 100 ml. of stock indicator were compared after successive additions of 0.05 N alkali to bring its pH up to 6.4. When correctly adjusted the color of the indicator solution is not changed by the addition of 0.2 ml. of the 0.001 ill buffer of the same rated pH, or pH 6.4 in this case. By this stepwise adjustment of the pH values of the several 100-ml. portions of the 0.001 ill stock indicator solution and the use of the same with the 0.001 M buffers to form the corresponding accurate “micro” buffer-color standard, great precision is obtained. Experiments were performed to see whether the stepwise adjustment can be omitted and the 0.001 M micro standards, prepared directly from 0.001 M indicators and buffers 0.4 pH or more apart, can be used to standardize the stock indicators accurately by use of 0.05 N acid or alkali from Table I of the former article (8). Trial adjustment from pH 6.7 to 6.5 actually gave pH 6.5; from p H 6.7 to 6.3 gave pH 6.3; and from pH 6.7 to 6.1 gave pH 6.15. Adjustment of the stock indicator can be made from values 0.4, or perhaps 0.5, pH higher or lower than the value desired, but not accurately from indicator solutions having pH values further removed, unless more concentrated buffers such as 0.05 M are used and corrections are made for their salt errors. The small dilution of the indicator occurring during the adjustment with the 0.05 N alkali is without practical significance, but may be compensated for by using an increased volume in high precision work. The volumes of 0.05 N alkali given earlier, Table I ( 8 ) ,for each pH value of the different indicators may be used for the preliminary adjustments of commercial indicators of varying purity. The final adjustments and frequent checks of these stock indicators should be made by the isohydric method described above. The indicator solutions are then ready for use. Nole-A set of 0.001 J I bromothymol blue solutions in Pyrex bottles equipped with small rubber bulbs on ground, hollow Pyrex stoppers sealed t o graduated capillary tubes for dropping t h e indicator showed less than 0.1 p H change with moderate use for one year. Soft glass bottles caused changes of 0.5 p H in one week. These 0.001 M indicators niay be stabilized against t h e action of t h e atmospheric carbon dioxide b y thorough aeration of t h e indicator solution during t h e p H adjustments or b y t h e establishment of t h e pH-carbon dioxide-bicarbonatwarbonnte equilibrium a s described in t h e first paper of this series ( 8 ) .
Khen the 0.001 .Vf buffers are used, especially above pH 7.0, the test tubes and other containing apparatus should be filled with Cos-free air and the tests made quickly to prevent pH changes. The 0.001 11f buffer-color standards should be sterilized and sealed in Pyrex test tubes, and when so treated are practically free from all errors. The data in Table I of this article indicate that the usual 0.05 M buffer-color standards have a salt error of +O.% to $0.3 pH ; that is, they have the same color or same colorimetric pH as 0.001 L’buffer-color standards having electrometric pH values 0.25 to 0.3 pH higher. Although the hydrogen electrode is itself probably subject t o similar and unknown salt errors, it is a t present accepted at; a standard means of comparing pH values of solutions. Any difference between the electrometric pH values of t’wo solutions showing the same uncorrected colorimetric pH is attributed generally, t’hough not entirely correctly, t o a salt error in the indicator method. As the 0.001 .I1 buffer-color standards are so dilute as t o be substantially free from salt effects, as shown below, the salt error is said to come from the more
80
A,VdLYTICAL EDITION
concentrated 0.05 M buffer and to increase with rise in the concentration of such salts and ions. K h e n the 0.05 JT buffer-color standards are used for obtaining the colorimetric p H of a very dilute solution free of salt errors, 0.25 to 0.3 p H must be added to the electrometric p H of the matched 0.05 M buffer-color standard to get the electrometric p H of the dilute solution. Likewise, when 0.001 JT buffer-color standards are used, 0.25 to 0.3 p H must be subtracted from their electrometric pH values to get the electrometric pH of a matched unknown of 0.05 N concentration. Table I shows that this rule holds good for the buffers used between p H 4.5 and 8.3. It is suggested that the 0.001 ill isohydric buffer-color mixtures be used as standards for measuring both the electrometric and the colorimetric p H values of unknowns. The salt correction will then be constantly zero for extremely dilute solutions; it must always be negative to gire the electrometric p H of the stronger solutions, the magnitude thereof increasing with the concentration up t o about 0.45 p H for 1 M phosphates. If 0.05 Jl buffer-color standards are used. as at present, the salt correction must be called plus (+) or added to the colorimetric pH to get the electrometric p H of more dilute solutions, and subtracted, or called minus (-), for more concentrated solutions. This confusion is ai-oided with the use of the 0.001 M buffer-color standards. The second and third coluinns of Table I taken together show the errors that arise if 0.2 cc. of the indicator, such as p H 6.1 bromothymol blue, is mixed with only 0.05 cc. (instead of 0.2 cc.) of the next p H step of the buffer such as p H 6.3, and the stock indicator is adjusted to this buffercolor standard and used as pH 6.3 without adding more alkali for the final adjustment as discussed heretofore. This procedure causes the p H of the adjusted indicator to fall just under p H 6.3, the error being increased at each step of the adjustment of the stock indicators. Hence the 0.001 “micro” standards made with 0.2 cc. of indicator and only 0.05 cc. of 0.001 JI buffer give the low colorimetric pH 1 alues shonm in column 3, instead of the correct ones shonn in column 2 , when compared with 0.05 JT buffer-color standards. The strictly isohydric technic described heretofore should be followed if the indicators are to be adjusted accurately in p H value. Table 11-Colorimetric pH of 0.001 M Buffer-Color Standards Exposed t o Carbon Dioxide of the Atmosphere for Various L e n g t h s of T i m e as Compared with 0.2 m l . Indicator Plus 10 m l . 0.05 M Buffera NO ExposrrRr: 2 HOCRS
4.3 5 06 6.1 6.25 6.45
6.65
,
.05 8.1
4 3 5 05 6 1 6 25 6.45 6 65 7 05 8 0
5 HOVRS
4 3 5 05 6.1 6.25 6 45 6.65 7.05 8 0-
1 DAY
4.3 5 05 6 05 6 22 6 42 6.62 7 00 7 8
3D~>s 4.3 5.05 6.06.276.46.55 6.95 7.4
7
DAYS
4 3
5.00 6.0f 6.2C 6.46.55 6.90 7.25
a 0 05 M buffers having the same electrometric pH \slues as those given in column 1 of Table 1 for 0 001 AI buffers.
Table I1 shows the gradual reduction of the p H of the more alkaline 0.001 Jl buffer-color standards when fully exposed to the carbon dioxide of the air. Changes in pH arising from moderate exposure in practical use may be neglected. As the 0.05 M buffer-color standards are very stable, it is perhaps better to use them in general practice and make the proper corrections for their salt errors. They may be specially adjusted and labeled to compeiisate for all positive or negative salt errors when used as standards in factories or research laboratories dealing constantly with very concentrated or extremely dilute solutions. The 0.001 iM buffer-color standards can be sterilized, protected from carbon dioxide, sealed in Pyrex, and used in special researches when salt errors are undesirable.
Vol. 2, No. 1
Technic for Using Adjusted Indicators
-1small drop of the unknown is mixed with a drop of each indicator as usual to find one or two indicators responding to the p H of the unknon-n. From 0.2 to 10 ml. of the unknown solution are added to 0.2 ml. of the proper indicator solution adjusted at the midpoint OF its useful pH range, buch as pH 7.0 for bromothymol blue. The comparison of this solution against the 0.001 M (or 0.05 31) buffer-color pH standards w t h corresponding volumes gives an approyiinate pH reading for the uiiknonm which is belom-, at, or above the midpoint p H of the indicator. Experience and judgnient make it possible to use the “approximate pH“ to formulate an “estimated pH” of the unknown. By mixing another sample of the unknown with 0.2 ml. of that indicator solution haxing the estimated pH and matching it against the buff er-color standards, there will usually be exact or nearly exact agreement between the p H of the indicator used and the pH as read; thus a new correct, or nearly correct, p H reading is obtained. Table 111-Comparison
of pH Values Found by Isohydric Method a n d by Other Methods INDICATOR PHDS. OBSERVEDPH
SOLETIOX ~IEASERED
CORRECT LOWEST POINT
MIDPOIST
HIGHEST POINT
Ind. Obsd. Ind. Obsd. Ind. Obsd.
P H BY SOHYDRlC
METHOD
Bromothymol blue
1 500 dilution of I C Lpeptone-lro glucose p H 6 5 medium
6 2
6 4 7 0 6 9 7 0 Bromocreiol p u r p l e
7.3
fi.9
1.1000 dilution of 1% peptone-beef infusion p H 6.05 bouillon
3.2
6 0
1. 1000 dilution of 0.03 .lilp H 7.82 phosphate buffer
5 8 6 0 6 0 8 6 6.3 Phenol red or creiol red
7 0
7 5 7.6 i.63 8 6 Bromocresol purple
1:5000 dilution of 0.02 .\I p H 4 . 2 1 phthalate bnber
5 0
1 10,000 dilution of 0 05 .If p H i . 8 2 phosphate buffer
7 0
7 . 2 7 . 4 7 . 4 7.8 Bromocresol purple
Air-CO? equilibrium p H 5.75 water
5 2
j 45
Copfree p H 7.0 superpure water
6.2
6.4
7 9
5 2 5 2 . 8 5 6 ti6 6 1 Phenol red or cresol red
7 05
3 ,4
7.65
7.4
0.0
5.75
7 . 0 7 05 7 . 6 7 2
7.0s
5.8 5 7 5 6 . 6 Bromothymol blue
K h e n the unknonn is scarce or expensive, 0.2 nil. may be riiixcd with 0.2 ml. of indicator and compared n i t h the “micro” buffer-color standards discussed above; the mixture of unknown and indicator, thus giving the approximate pH, can be mixed with another 0.2 ml. OF that stock indicator solution selected from Table I of the first article of this series (8)to make the total indicator isohydric with the solution with the estimated pH. Comparison JFith special “micro“ buffer-color standards having double quantities (0.4 ml.) of indicator will then give good pH readings for the unknown. In general, barring salt errors discussed below, correct pH values may be obtained by using indicator solutions adjusted (1) a t their midpoint pH for measuring the pH valueq of the usual buffers and media and up to about 50-Fold dilutions thereof; ( 2 ) a t their lowest, midpoint, and higheit useful p H values for studying such buffers and media diluted about 50 to 1000 fold; and (3) in 0.2 pH steps over their useful ranges for measuring the p H values of very dilute buffers, media, and water. This isohydric indicator method gives a new technic for measuring the pH of solutions niorc dilute than 0.0001 -11 with which the electrometric methodi are either inaccurate or entirely inapplicable. Tables 111 to 1-11show the inagnitude of the errors found when pH
Ih-DCXTRIAL AA7D ENGINEERING C H E N I S T R Y
January 15, 1930
measurements are made on very dilute solutions or water without using the isohydric indicator method. Table I11 shows a few p H measurements on media, buffers, and water which illustrate the application of the isohydric indicator technic. The lowest useful p H values of the indicator and the observed pH’s of the mixture of this indicator and the buffer solution are given first. This lowest point adjustment of the indicator requires about one molecular equivalent of alkali and is the procedure employed usually without any consideration of the p H value. The column headed “midpoint” shows whether the approximate and estimated pH’s of the solution are below, at, or above the midpoint pH, and leads to the final check of the p H of the unknown with the isohydric indicator as given in the last column. The observed p H values for the highest and lowest points are to be compared with those in the last column and are especially helpful in showing the errors due to the use of indicator solutions that are either unadjusted or adjusted a t a p H value quite different from the p H of the unknown, weakly buffered solution. This isohydric indicator technic is also very useful in the measurement of the colorimetric p H of weakly buffered solutions containing large concentrations of neutral salts and having large salt errors. The electrometric reading given by a hydrogen electrode in such a solution is ordinarily used as a measure of the p H thereof. It is uncertain, however, whether the electrode and the e. m. f. readings are affected by the salt, since large amounts of such salts sometimes render the electrode inoperative. But in practical work with concentrated solutions of constant content of neutral salt and fixed though unknown salt errors, but variable low buffer value, the isohydric indicator technic gives the colorimetric p H quickly and accurately. The salt-error corrections can be applied later when they are determined by proper electrometric tests. Production, Test, Storage, and Use of Superpure Water for Precision pH Work The solid contents and carbon dioxide found in commercial once-distilled waters having p H values generally of about 5.0 to 5.5 and sometimes up to 6.0, render them unsuitable for precision p H work. The writers have therefore begun a study of the effects of varying amounts of these impurities T a h l e V-Dilution
I
PHOF INDICATOR
6.2 6.4 6.6 6.8 7.0 7.2
7.4 7.6 6.2 6.4 6.6 6.8
7.0 7.2 7.4 7.6
II
PH OF UNDILUTED BUFFER ElectroCoiorimetric metric
7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00 7.00
7.0
5
7.0 7 0
I\LIII”I
6 7
2.p
I t is the writers’ experience that for p H work it is far more important to try to remove all of the carbon dioxide than all of the solid neutral electrolytes and both conductivity and p H tests of the purity are necessary. For example, water ha\-ing the specific conductivity 0.8 X a t 18”C. might molar concentration of the non-buffer, contain a 6 X
7.15
7.0+ 7 . 0 ~ 7.0t 7.0+ 7.0+
7.15 7.15
7.15 7.15
7.15 7.15 7.15 2.16
7.15 7.15 7.15
7.16 7.15 7.15
7.15 7.15 7.15 7.15 7.15 7.15 7.15 7.15
7.15 7.15 7.15
,,15
2.15 r.15 7.15 7,l.i 7.15
OF
DILUTED 11EDIUM
7.10
7.02
6.95 , 00 7.02
7.0+ 7.0+
7.0f 7.0+
PH
7,l5 7.15 7 . i5 7.15 7.15 7.15
7.2
7.15
cu
( 2 ) Fresh soda-lime in jar without Chamberlain filter to catch soda-lime dust (0) (b) (C) 0 7.06.85 7.1 ’/? 7.07.0 7.l+ .. 1 7.2 2 1.5 7.5 7.4 3 .. .. 7.6
1:10,000
7.2 7.2 7.2 7.2 7.2 7.2
7.15 7.15
7.1 7.05, 7.07
7.05 7,0+
1:1000
7.15
7.15
7.0 i. OL,
1:500
7.15 7,l5
7. 1.5
7.0-t
7.07 0-
1:100
7.15
7.01-
1 2 3
1:50
7.15 7.15 7.15 7.15 7.15 7.15
p H 7.05 with bromothymol blue (7.0). b p H 5.7 with bromocresol purple ( 5 . 7 ) . c Direct from still.
a
1:10
7. 1.5
7.0+
T a b l e IV-pH Values of S o d a - L i m e S c r u b b e r s and S u p pure Water .e r. (1) Chamberlain filter between used soda-lime jar and Basks 1 and 2; flask 3 attached directly to jar
of F e r m i ’ s S o l u t i o n B r o m o t h y m o l B l u e
7.0+
7.0+ 7.0+ 7.0+ 7.0+
and of the best methods for their removal to produce highgrade water for practical use in the laboratory. The researches of Kohlrausch ( I O ) , Bourdillon (4) Kraus (IL), Washburn (153, Wieland (IC),Kendall (9). and others have s h o m that repeated distillation of water in a stream of C02-free air or steam or in a vacuum will produce what n e shall call “practically absolutely pure” water free from extraneous electrolytes and having a specific conductivity of about 0.04 X a t 18” C. Such water has been prepared in small volumes as needed, and it is very difficult to preserve and use it without special t.quipment. Since their nork has s h o m that for even precise colorimetric p H work there is a limit beyond which it is futile to purify the water, the writers have concentrated their efforts on the production, test, storage, and use of a double-distilled water that meets the requirements and can be prepared efficiently without special attention. Such water occupies the position bespecific tween absolutely pure water of about 0.04 X conductivity and the best of the conductivity water used before 1910 with a specific conductivity of aboilt 0.8 X a t 18’ C. It is often called “superpure” water, but will be designated also as p H 7.0 water in this article.
COLORIMETRIC
1:5
81
7. .~2
7,l5 i,l5 i,l5 7.15 7,l5 7,l5 7,l5 7.15
7,l5 7,l5 7.15 7.15 7.17 7 15
2. 15
1.15
7.15 i,l5
7,l5
7.15
7.15 7.15 7.15 7.15 7.15 7.15 7.15 7.15
7,05
6.85
7.15
7.00
7.05 7.05 7.05 7.05 7.05 7.05 7.05 7.10
6.5
6.0
i.05 7.10 7 15 7.17 7.20 6 1 6.1 6.2 6.25 6.3
6.4
6.5 6.6
1:100,000 6.6 6 7
6.75 6 85 7.05 7.2 7.47.4
CO>-Free Water0
6.4
6 5 4
6.75 6.85 i.05 7.?+ 7.47.5 Air..COr Equilibrium Water b 6.06.06.0 6.0 6.1 6.2 6.3+ 6.35 p H 5.0 Waterc
F~illacid color 6.6
6.2
A X A LY TICAL EDITION
82
neutral potassium chloride giving a pH reading close to 7.0. It might, on the other hand, contain 14 X mol per liter of carbonic acid (Cos H~COS), giving it high buffer properties and a p H of 5.7. I n other words, the quantity of carbon dioxide dissolved in pure water a t equilibrium with atmospheric carbon dioxide and giving a pH of 5.7 is sufficient to account for the conductivity usually found for what is known as "good conductivity water." It is therefore clear that practically all of the errors in pH data arising from the use of good conductivity water are due to carbon dioxide, even though it is granted that a small percentage of the residual solids such as carbonates have some buffer properties and may give the p H 7.1 to 7.3 or higher in the water when all the carbon dioxide is removed. All the more readily, then, can serious pH errors arise by the use of ordinary once-dis-
Vol. 2,
KO.
1
thymol blue of pH 7.0, as shown in Table IV. The pH 7.0 water used in this work was always of this quality. Even with a correction of 0.2 to 0.3 pH for the salt error of the 0.05 M buffer-color standard, the pH is less than 7.4. The insignificance of a variation of *0.2 p H in practically pure water may be judged by the fact that it would take only 1.46 X IO+ equivalent of completely ionized carbonic acid, hydrochloric acid, or sodium hydroxide to make such a change. Such a trace would produce a change of only 0.02 pH in a solution of hydrochloric acid or sodium hydroxide a t pH 6.0 or 8.0, and would be without any measurable effect on any commonly used buffer solutions. But even this small amount of impurity is detected by the isohydric indicator standards, though not with the usual indicators. This impurity expressed as carbon dioxide is only about 1 per cent of the amount needed to give the specific conductivity 0.8 X IOp6 to so-called good conductivity water and hence within limit of measurement thereof. The amount of solid obtained by evaporating 30 liters in a quartz flask and platinum dish was never more than 0.15 nig. per liter, and was lowered to 0.03 to 0.06 mg. per liter by frequent use of dilute acetic acid and pH 7.0 water between distillations to wash accumulated solids from the stillhead A . If this solid were all potassium chloride, the specific conductivity should be less than 0.25 x 10-6 a t 18" C. One conductivity test on the water under none too rigorous conditions gave a specific conductivity Iess than 0.25 X 10-0 at 18" C. It is believed that such a grade of superpure water is suitable for the preparation and tests of extremely dilute, 1% eakly buffered solutions with isohydric indicators to within 0.05 pH and further distillations will then give practically Figure I-Table Still a n d S o d a - L i m e Scrubber pure water when higher precision is needed. tilled water having about a 300 X molar concentration S. F. Acree and Fred Acree, Jr. ( 1 ) have announced the deof carbon dioxide and a pH of 5.0 in addition sometimes to velopment of a continuous commercial still embodying these sufficient total buffer solids to give pH 8.0 and higher when principles and capable of producing similar superpure pH all carbon dioxide is removed. 7.0 water containing less than 0.1 mg. solids per liter and A table still as shown in Figure 1 has therefore been used having a specific conductivity below 0.3 X IO+ a t 18" C. to redistil the usual grade of distilled water and collect as With such water once prepared, it is easy to pass fresh little as possible of the solids and carbon dioxide in the dis- air through it to get pH 5.7 air-CO? equilibrium water or to tillate. This general type of apparatus has been shown by put in carbonated water to give it any desired reaction such other workers cited above to give practically absolutely pure as pH 5.0 found in many ordinary distilled waters. This water by repeated distillations. The writers have added a has been done, and buffers and media have been diluted block tin or Pyrex stillhead, A , filled with tin or Pyrex rings, therewith and studied by both the electrometric and isohyand have connected it with a rubber stopper covered with dric indicator technic to measure the pH changes and salt tin-foil, or with a ground joint, to the 5-liter Pyrex flask B. and protein errors. The results are given in the following The condenser tube, C, is of block tin. Compressed air is sections. passed downward in the jar, D, through moist soda lime and Superpure water shouId be stored in clean Pyrex or blockthen upward and out through a Chamberlain filter (not tin reservoirs in an atmosphere of COz-free air obtained from shown) in a central tube, E . This modification of the well- the soda-lime scrubber. In making up the solutions for known Henderson-Carpenter apparatus removes both the pH tests such water should be transferred by siphoning into carbon dioxide and all soda-lime dust from the air stream. the graduated flasks or other equipment after they are filled The Chamberlain filter may be used outside the soda-lime with COz-free air. It can also be transferred safely and jar, but omission thereof or the use of certain less porous readily into the test indicator solutions or the buffer filters studied extensively can easily cause the water to reach solutions by means of a Pyrex pipet, provided the tip dips pH 8.0 and above from soda-lime dust. This soda-lime-filter under the liquids and pH readings are made quickly. air-scrubber is very convenient for preparing pure COzTable I\' shows some data on the pH of the water obtained free air and with moderate use will last a year with only from the table still, on the efficiency of the soda-lime scrubber one filling with soda lime. When such COz-free air is passed in removing carbon dioxide from the air, and on the value through pH 5.0 distilled water heated in the flask B , the of the Chamberlain filter in separating fine soda-lime dust carbon dioxide is quickly removed, and the first quarter from the air stream. Three flasks of redistilled water were of the distillate rarely has a pH below 6.9 when tested with collected as fractions 1, 2, and 3 from the table still, all three bromothymol blue of pH 7.0. samples testing about pH 7.0 with pH 7.0 bromothymol blue. The entire distillate is condensed in one stream of ConAir was passed consecutively through a jar, D, containing free air and collected in another one in the Pyrex flask F . used soda lime, a Chamberlain filter, and flasks containing Both air streams pass out a t G. The stillhead, A , prevents fractions 1 and 2. A flask containing fraction 3 was attached any appreciable spray containing buffer solids from passing directly to the soda-lime jar. Table IV (1) shows that the p H into the condensate. When such a distillate testing pH did not change in one or more weeks, and hence that no appre7.0 or 7.05 is aerated for weeks with the Cos-free air, the ciable amounts of carbon dioxide or soda-lime dust came pH nerer rises above 7.05 or 7.1 when tested with bromo- through the jar of used soda lime. On the other hand,
+
I S D C S T R I A L All-D EXGlNEERIA-G CHE-TIISTRY
January 15, 1930
83
of 1 M P o t a s s i u m Acid P h o s p h a t e - N a O H Buffer w i t h COz-Free ( p H 7.0) W a t e r
T a b l e VI-Dilution
SAL,T ERRORS
COLORIMETRIC PH DILUTION PARTS
lIOLALITY
ELECTROMETRIC PH
0:o 1:s
1.0 0.2
7.80 7.86
1:lO
0.1
7.89
1:20 1 : 100 1 : 500
0.05 0.01 0,002
7.9: 8.01 8.12
1 : 1,000 1 : 5,000 1 : 10,000
0.001 0.0002 0.0001
8 16 8.12
P H E N O L R E D ADJUSTED TO:
7 . 9 2 av.5
I
8.0 p H
8.08.0-
7.95 8.07.98b 8.07.98b 8.05 7.98'' 7.97t
8.08.0-
7.95 7.95
7.9 i.8 7.65
7.9b 7.85;n 7.805
8.07.85 7.85
7.75 7.6 7.4-
XOLALITP
1:10,000 1: 50,000 1 : 100,000 1 : 500,000
0.0001 0 . 00002 0.00001 0.000002 H10
8.6 p H
~
PH
WITH
7 . 2 pH
7.4 pH
7.6 p H
7.8 pH
8.0 p H
7.65
(7.7-) 7 55 7.4 7.25 7 2
(1.7)
7.75 7.67 7.55 7.6 7.5
7.8 7.75 7.65 (7.65) 7.65
7.2
-~
,.62 7.43 7.4 7.4-
Ioshydric solns.
Alcoholic solns.
7.95 8.07.986 7.95b
+0.1:1 +0.12!
8.08.08.0!.95h .9b 7.851
+0.03 -0 03 -0.1Cl
-0.06 -0.17
-0 2CI -0 27 -0.28
-0.41 -0.60 -0.60
+0.09
,
7.755
~-
COMPLETE SERIESOF PHEXOL RED
7.0 p H
7.0
8.0 p H
Alc. soln.
Isohydric Serial Method, Con-Free W'aterc COLORIMETRIC
DILUTIOS
To
7.0 pH
i
PHENOL RED
CRESOL R E D
~
-1
8.2 p H
8.4 pH
8.6 p H
7.8
7.85
7.85
7 . 77 8 (7.7) 773
z 5 7.8 7 8
(1 82)
7.9 7.85 7 CJ 7 0 7 9.5
,
85 785 7 !2
ISOHYORIC PH
7.3
a Average of seven readings ranging from pH 7.8 t o 8.02. This columri and other d a t a show t h a t in concentrations of 0.0001 M and below t h e hydrogen electrode is unreliable and may give readings varying T h e isohydric indicator method, however... gives t h e colorimetric D H of ~.several tenths of a D H unit. such solutions within 0.1 unit. b As t h e p H 8.0 indicator gave readings distinctly below pH 8.0 in some cases these measurements were estimated by the isohydric method. c Continuation of first section of table; solutions too dilute t o measure with electrode. ~
when a jar of fresh soda lime was attached directly to three other fractions, a, b, and c, the air stream transferred dust particles from the fresh soda lime into the water and raised the pH from about 7.0 to 7.5, as shown in Table IV (2). I n some cases pH 8.0 has been reached. Although long use of a soda-lime jar at times eliminates dust from the air stream, absolute safety is obtained only by the use of the Chamberlain filter. pH Changes Caused by Dilution of Fermi's Solution with Superpure, Air-Ewilibrium, and Ordinary Distilled Water
This is not surprising in riew of the fact that pH 5.0 water contains about 33 X mol of total carbon dioxide per liter, pH 5.7 water has 1.4 X 10-5 mol per liter. and pH 7.0 water practically no carbon dioxide. The effects of the buffer salts in the medium on the pH are to be seen in dilutions as high as 1:100.000. The phosphate concentration is then only 8.3 X 10-7 Jf, but e\-en this .lightly changes the pH of the combined indicator and water. It is especial57 noticeable with the pH 7.0 water and slightly so with the p~ 5.7 water, ~~t at a dilution of only about 13250 with the pH 5.0 water the carbon dioxide is equal to the phosphate in concentration; as its pH (5.0) is lower than that of the phosphate (pH 7.0) and as carbonic and secondary phosphoric (H2POI-) acids have practically the same ionization constants, 3 x 10-7 and 2 X lo-', respectively, the higher hydrogen-ion concentration due to the carbon dioxide causes the marked divergence of 0.40 pH from the other curves a t this dilution. This increases to as much as a pH unit a t 1:lOOO dilution. Beyond this point the pH is below 6.0 and therefore below the range of bromo-
Table V and Figure 2 show the p H errors that may be expected when a well-buffered medium, such as Fermi's solution containing phosphates, magnesium sulfate, and glycerol, is diluted with (1) double-distilled Coyfree pH 7.0 water, ( 2 ) the same water a t air equilibrium (pH 5.7), and (3) ordinary distilled water taken directly from a commercial still and having the p H 5.0. These pH readings of the waters were made with the suitable series of adjusted indicators. The Fermi's solution used was standard (8) H , M and contained 0.083 M phosphates. Its electrometric pH was 7.0. The table and curves illustrate well the fact that on dilution a medium may a t first show buffer-dilution changes in pH arising partly from M increased ionization and hydrolysis of the buffer salt and also partly from the salt effect or salt 90 errors due to the effect of varying concentration of the buffer salt on the color of the indicator a t constant electrometric pH value. The differentiaM tion of these factors by the use of both electrometric and isohydric indicator methods is now 70 being investigated in detail. The increase of 0.15 pH unit first observed is in the opposite di72 rection from the decrease in pH or higher acidity produced by the carbon dioxide when the medium I is diluted beyond 1:lOO with pH 5.7 water and t I *.*r with pH "O water' This dioxide * z u r e 2-Dilution of F e r m i ' s S o l u t i o n . B r o m o t h y m o l Blue i n A d j u s t e d Series has an increasing influence as the medium is di( p H 6.2-7.6 i n I n t e r v a l s of 0.2 p H ) luted further, and it is clearly seen that three distinct zones of pH values are produced by the wide varia- thymol blue. Any further dilution with the pH 5.0 water tion in the amounts of carbon dioxide present in any would naturally approach pH 5.0 and require the use of given dilution with p H 5.0, pH 5.7, and pH 7.0 water. bromocresol purple and bromocresol green. It should be HW
HMO
k1000
110000
1 9 0
A X 8 L Y TICAL EDITION
84
noted especially that the graded adjustment of the bromothymol blue from 6.2 to 7.6 causes a divergence in the corresponding pH readings in the higher dilutions beginning with about 1:jOO. The table and curves show clearly that the pH 7.0 Fermi’s solution gives constant readings t o within 1 0 . 1 5 pH a t all dilutions with the isohydric pH i.0 water and pH 7.0 indicator, but waters and indicators of any other pH values produce noticeable variations from these correct isohydric readings. T a b l e VII-Salt Electrometric
INDIcaroR
Phenol red
Bromothymol blue
Bromothsmol blue
a n d P r o t e i n Errors for S u l f o n p h t h a l e i n I n d i c a t o r s . I s o h y d r i c C o l o r i m e t r i c D a t a , Using 0.05 M BufferColor S t a n d a r d s
YS.
C O N C N . A N D PH OF L I E D I E ? OR BUFFER
Phosphate-NaOH, p H 7.60 1.0 mol 0 2 mol 0.1 mol 0 . 0 5 mol 0 . 0 1 mol 0.002 mol 0.001 mol 0,0002 mol Phosphate-NaOH, p H 6.80 0 . 0 5 mol 0.005 mol 0.001 mol 0,0005 mol 0 , 0 0 0 1 mol 17 Peptone-1% glucose groth p H 6.88 (weaker t h a n b.05 M buffer-color standard) Undiluted 1:lO 1:20
SALTERROR ’ROTEIN
0.001 ;\I 0.05 *\I stds. stds.
P 11 +0.43 +0.35 +0 30 +0.26 +0.17 +0.06 0.00 0.00
+0.15 t0.12 +0.09 A0.03 -0.03 -0.15 -0.26 -0.25
+0.20 +0.05 0.00 0.00 0.00
0.00 -0.15 -0 20 -0.20 -0.21
-0.13 -0.22 -0.23 -0.23 -0.23
1:50 Bromocresol purple
Bromocresol purple
1:100 Phthalate-h-aOH, p H 6.01 0.05 mol 0.01 mol 0 001 mol 0 0005 mol 0 , 0 0 0 2 mol Peptone-beef infusion, p H 6.22 (about same concn. a s 0.05 M buffer-color standard) Undiluted 1.. l o 1:50 1:lOO 1:500 Phthalate-NaOH, p H 5.22 0.05 mol 0.005 mol 0.001 mol 0.0005 mol 0,0001 mol 4-Day B. coli culture in 170peptone-l% glucose broth, p H 4.78 Undiluted 1:s
+0.23
+0.12 0.00 0.00 0.00
0.00 -0.11 -0.23 -0.23 -0.23
-0.04 -0.17
~~
Bromocresol green
Bromocresol green
-0.20 -0.23
-0.23 L O 20 fO.05 0 00
0.00 0.00
0 00 -0.15 - 0 20 -0.20 - 0 20
-0.18 -0.15
1:10 150
Bromophenol blue
Bromophenol blue
1:100 1:500 Phthalate-NaOH, pH 4.22 0.05 mol 0.01 mol 0.005 mol 0,001 mol 0.0005 mol Peptone-beef infusion 20 cc. 1 N HCI per liter, p H 4.16 (stronger t h a n 0.0: M buffer-color standard) Undiluted 1:5 1:lO 1x50 1:100 1500
+
ERROR
-0.14 -0.18
-0,l5 -0.16
t0.16 t0.07 t0.03 0.00 0.00
0.00 -0.09
-0.13 -0.16
-0.15
C O . 15 0.00 0.00 -0.10 -0.14 -0.12
Some work has been done on bubbling air through various calculated mixtures of sodium bicarbonate and carbonate t o get known (302-air mixtures, which are then passed into superpure water to give it the desired pH value due to the carbon dioxide. This suggests that water can be made up a t any desired pH by this method or by the use of carbonates, ammonium salts, sulfites, sulfides, hypochlorites,
VOl. 2, No. 1
etc. Such waters can then be used for these dilution studies, Relation of Dilution to Salt and Protein Errors
In the preceding sections methods were described for measuring the correct isohydric colorimetric pH values of buffers diluted with superpure water; but the colorimetric pH reading is progressively changed by decreasing the concentration of the buffer solution while keeping it a t the same electrometric pH. The electrometric pH value is chosen as the standard and the deviation of the colorimetric pH therefrom is called a “protein error” when protein buffers or media are used and simply a “salt error” when other organic or inorganic buffers are used. The isohydric colorimetric pH yalues differ from the standard electrometric pH values by the amount of salt or protein error. Table 1‘1 and Figure 3 show data for a common 0.05 JI phosphate buffer of given electrometric pH mixed with an indicator to form a buffer-color standard whose colorimetric pH is usually taken as equal to the electrometric pH for practical use. On dilution of this 0.05 M buffer the isohydric colorimetric pH values obtained by comparisons with the 0.05 Jf buffer-color set gradually became smaller than the corresponding electrometric pH values until a dilution of about 0.001 Jf was reached. Further dilution gave a nearly constant difference of 0.25 to 0.3 pH. On the other hand, a gradual increase in the concentration of this buffer up to 1M caused the series of pH values determined colorimetrically against the standard 0.05 M buffer-color set to become steadily larger than the electrometric pH values a t the same concentrations. Since the choice of 0.05 ill standards, or the intersection of curves A and B , is purely arbitrary, it is believed that the correct interpretation of the data is the assumption that at a dilution of 1:10,000 the sinall concentration of buffer (0.0001 M ) produces no salt error. The apparent colorimetric pH curve, B , should therefore be moved upward to take the position of the true colorimetric pH curve, C, which is free from salt error. The concentration of the indicator is constant and preferably 0.00002 AI and hence produces no salt effect, It is then seen that the increase in the buffer concentration above 0.0001 M causes the isoactive colorimetric p H values of curve C to remain approximately constant while actually diverging steadily from the electrometric pH readings, the real salt errors being ah-ays positive and the corrections negative. Similar data on the salt and protein errors for a number of the sulfonphthaleins and methyl red when mixed with the commonly used buffers and media in widely varying concentrations in superpure pH 7.0 water have already been found. All unknowns should be thus measured both colorimetrically and electrometrically when salt or protein errors are of importance. Columns 7 and 10 in Table VI show the errors that may be expected when alcoholic solutions of the unneutralized indicator are used. The use of ordinary distilled pH 5.0 water, or even air-COz equilibrium pH 5.7 Rater, and unadjusted indicators should be avoided as it leads t o variable and erroneous values for the colorimetric and electrometric pH and the salt and protein errors in the very dilute solutions. Salt and Protein Errors for Some Common Indicators, Buffers, and Media
The salt and protein errors (see 3, for study of salt errors) recorded in Table VI1 were calculated against both 0.001 and 0.05 M buffer-color standards from data obtained by the use of superpure water, isohydric indicators, and electrometric measurements as discussed in the preceding section. The buffers and indicators studied covered a p H
I S D C X T R I A L A S D ESGINEERIA-G CHEMISTRY
January 1.5, 1930
range from 3.0 to 8.6, which meets most needs. The dilution of the buffers and culture media ranged from 5 to 5000 fold. The data for thesalt errors against 0.001 Mbuffer-color standard\ s1ion.n in column 3 of Table VI1 were obtained as follows: The pH values of each original solution and the dilutions thereof were obtained first electrometrically by means of the Acree double hydrogen electrode, and then colorimetrically by the isohydric method against the 0.0.5 J1 buffercolor italiclards and in some cases against the 0.001 -lI series as chechb. The colorimetric pH values estimated from the 0.05 -11 buffer-color standards were plotted to give a smooth curl-e for each dilution series; a corresponding smooth curve n a5 made for the electrometric pH values. The tTTo curves were found t o shoTT deviations called the “salt error“ and t o reach a constant difference in pH when the dilution was greater than 1:lOOO. This difference varies with the dilution, the buffer, and the indicator, as shown in column 4 but is of the order of 0.2 to 0.3 p H unit. A new colorimetric pH cur\ e mi..then drawn parallel to the first one and coincid-
85
tained against 0.05 hl buffer-color standards are therefore large and of the order found for the dilute standard buffersnamely, 0.15 to 0.25 pH units. On the other hand, the peptone-beef infusion is not only naturally well buffered, but when treated with sufficient 1 -2; hydrochloric acid to shorn pH 4.15 the increase in qalt and buffer content makes it behave on dilution more like 0.1 or 0.2 II phobphates. In other words, it has a +O.l5 pH salt or protcin error. which decreases to -0.12 pH 011 500-fold dilution. The same medium TT ithout any added hydrochloric arid belial es more like 0.05 31 phobphates in that the protein error changer from -0.04 pH to -0.23 pH on 500-fold dilution with the pH 7.0 water. I t must not be forgotten that these culture media contain both soluble buffers and buffer-like colloidal proteins, whose state of aggi egation and electrical charges and adsorptive poner for the iiidicatcir change with pH and with dilution and cause uncertainties in any iiitcrpretation concerning protein errors. d fundamental interpretation of the cause and mechanism of salt and protein errors, as well as of‘ the contact potentials invoked in the electrometric pH determinations, cannot be giT en for the more concentrated buffer solutions until more data are available on the viscosities, ionic and molecular mobilities, dielectric constants, transference numbers, hydration, activities as determined by various methods, and other physical-chemical constants in-rolved. Some notable advances along this direction have been made For the more dilute buffer solutions whose phyiical-chemiral I-onstants are better knonn (gj 5, e. 7 , 11, 13, I $ ) . Literature Cited
1
I
I
,
,
/
I
11 /
p
675
/ I i l j I MLyJ;,y1 1 1 1 I
~
/O
GI
OWOl
O
0 OWOI
Figure 3-Relation of Electrometric and Colorimetric pH Values to Dilution of Phosphate Buffer. Phenol Red a n d Cresol Red (pH 7.0-8.6)
ing with the electrometric pH curve at concentrations of about 0.001 to 0.0001 Jd. This new curve is therefore free of any salt errors arising from the 0.05 K buffer-color standards and gires the correct isohydric colorimetric pH values, and its deviation from the electrometric pH curve gives the true colorimetric errors or salt errors shown in column 3 of Table F’II. As an example, the data for phenol red and dilutions of 1 IId phosphates are taken from Figure 3. A11 the sulfonphthaleins give positive salt and protein errors or negative salt corrections of about the same general magnitude against 0.001 M buffer-color standards. The salt corrections against the usual 0.05 Jd buffer-color standards are generally negative for solutions stronger than 0.05 -11 and positive for solutions weaker than 0.05 M . The colorimetric pH values for the various culture media vere not plotted against 0.001 -11 buffer-color standards. The actual concentrations and identity of the buffers in such complex media are not known, and the pH curve against the 0.001 M standards, when needed, must be expressed in terms of dilution of the media instead of the molar concentration given for the organic and inorganic buffers of column 2. Severtheless certain generalizations may be given for the salt or protein errors of the culture media. For example, the 1 per cent peptone1 per cent glucose broth contains naturally less buffer, or has a smaller concentration of buffers. than the peptone-beef infusion. This is shown in their pH titration curves. The salt or protein errors ob-
(1) .%Cree a n d hcree, Paper presented b e f o r e t h e Division of T\-ater, Sewage, a n d Sanitation Chemistry at t h e 77th Meeting of t h e hnierican Chemical Society, Columbus, Ohio, .%pril 29 t o M a y 3, 1929. (2) Bjerrum, “Die Theorie der alkalimetrischen und azidimetrischen Titrierungen,” p. 1, 1914. (3) Blum and Bekkedahl, Tvans. .Am. ElecLrociiem. Soc., 56 (1929). preprint. ( 4 ) Bourdillon, J . Chenz. SOL.,103, 731 (1913). ( 5 ) Brdnsted, I b i d . , 119, 574 (1921). (6) De Bye, Z.p h y s i k . Chem., 130, 56 :1927), (7) D e B y e and BIcCauley, Physiii. Z.,26, 2 2 (1923). (8) Faivcett and Acree, J . B Q C ~17, . , 171 (1929). (9) Kendall, J . Chem. Soc., 101, 1275 (1912); J . A m . C h e w . Soc., 38, 1480
(1916).
(IO) Kohlrausch and Heydweiller, 1l.eid. . I n n . , 53, 209 (1594); Z.g h y s i k . Chem., 14, 317 (1894). (11) Kolthoff, J . P h y s . Chem., 32, 1820 (1928). (12) K r a u s and Dexter, I b i d . , 44, 2465 (1922). (13) Sorenson, Compl. v e n d . Iraa. l a b . C a u l s b e ~ y ,8, 1 (19Cl9). (14) SBrenson, 2 . Biochem., 21, 159 (1900). (15) R’ashburn, J . .Am. Chem. Soc., 40, 106, 1 2 2 , 150 (1918). (16) \vieland, I b i d . , 40, 131 (1918).
Sunchecking of Rubber For some years the rubber industry has been active in developing materials to retard deterioration in rubber. Seireral materials now employed are of decided advantage when the rubber is not exposed to sunlight, but a different type of material is demanded for protection from sunlight. The Bureau of Standards during the summer made tests on a commercial material designed to prevent or retard cracking and checking of rubber compounds exposed to the sun. Samples of rubber tubing were made containing antichecking material, and other samples with a wax commonly used for this purpose. With the tubing under a slight tension checking was apparent on plain samples within 15 days, while samples containing “anticheckers,” or large amounts of wax, showed none for a t least 90 days. Waxes in the amounts necessary to prevent sunchecking have a tendency to “bloom” to the surface-a feature which is objectionable 011 some rubber articles. Tests indicate that commercial antichecking materials might be used in many places where the surface bloom of wax would be objectionable.
