Rapid Electrometric Determination of the Alkalinity of Sea W a t e r Using a Glass Electrode DON H.ANDERSON1 WITH REX J. ROBINSON Chemical and Oceanographic Laboratories, University of Washington, Seattle, Wash.
O n the basis of glass electrode p H measurements of solutions to which known amounts of dilute hydrochloric acid have been added, “empirical coefficients” .have been calculated, and the normality of a dilute solution of sodium bicarbonate containing a mixture of salts has been determined with a precision of about 1%. .The operations are illustrated in connection with determination of the alkalinity of sea water. O n e measurement of p H yields the normality or alkalinity directly and a direct-reading pH instrument could b e scaled to read normality, molarity, or alkalinity directly. The factors that affect the accuracy of the determination are listed and their magnitudes indicated. All steps are indicated that must b e performed in developing a similar procedure for systems other than sea water. Industrial control determinations, and clinical investigations where large numbers of determinations are made, can b e rimilady standardized.
THE
term “alkalinity” has been adopted by the International Association of Physical Oceanography (1939) as the standard designation to replace the terms ‘(titratablebase”, “excess base”, “titration alkalinity”, and “buffer capacity” of sea water. It is defined as being the number of milliequivalents of hydrogen ion neutralized by 1 liter of sea water a t 20” C. The alkalinity of sea water is primarily due to the presence of bicarbonates, though borates, phosphates, arsenites, and silicates are present in lesser amounts to form a complex buffered system with a pH of about 8. Information concerning t’he total concentration of salts of these weak acids as well as the available alkali cannot be acquired from a single pH measurement but must be determined otherwise. The methods for the determination of alkalinity of sea water have been well summarized by Sverdrup, Johnson, and Fleming ( 3 ) . Because of t,he variety of results obtained by the different methods, there is a definite need for a standard procedure yielding accurate and precise results which can be universally understood and compared. Sone of the methods involves any assumptions as to the nature of t,he buffers present in sea xater, but all take into accouiit the variation in salt content. The work of RIitchell and Rakestraw ( 2 ) Tms based upon a titration procedure and yielded excellent results. In the method of Thompson and Bonnar ( 5 ) an excess of standard acid was added to the sea-water sample. The amount of the acid that had been effectively neutralized by the sea water \vas readily determined by measuring the amount of excess acid colorimetrically, by comparison with n series of standard tubes. Though the method i? convenient’, the stnndard comparison tubes n-ere difficult to prepare and rather unstable. Some time ago the glass electrode was used at this laboratory by Thompson and Anderson (4) for making routine pH and alkalinity measurements of sea water. ’ This eliminated some of the difficulties of the ‘earlier methods, but ne\\- considerations arose. Further investigation of the u.se of the glass elcctrode by the present authors in determining the concentration of the excess acid has demonstrated the necessity of a knowledge of
the relationship between the pH of the aolution and the amount of acid added to sea water of various salinities. This involves the experimental determination of coefficients which are empirical, having significance only when applied to this determination with the specified conditions. DETERMINATION OF EMPIRICAL COEFFICIENT
The empirical coefficients were determined from a series of titrations of natural and artificial sea-water samples. Sea water is alkaline to about the same extent as a 0.002 N sodium bicarbonate solution. One hundred milliliter samples were titrated with 0.01000 N hydrochloric acid until the pH was about 3.0. The equivalent point was taken as occurring with the greatest change in e.m.f. per unit volume of added acid, a t a pH of about 4.5. The pH of the solution was measured a t frequent intervals beyond the equivalence point. From the known amount of acid added and the experimentally determined alkalinity, it was possible to determine the excess acid present. The empirical coefficient, fa+, was calculated by dividing CH+ f ~ + obtained from the pH measurement of the solution by the concentration of excess strong acid, CH+, as shown under calculations. From this coefficient, CH+ in other similar solutions can be calculated. EXPERIMENTAL EQUIPMENT.The pH measurements n ere made with Leeds I%Northrup glass electrodes No. 1199-12 and reference electrodes KO.1199-13. These were connected with a Leeds Bt Korthrup vacuum tube amplifier to a Type K2 potentiometer. EXPERIMENTAL PROCEDURE.The electrodes were standardized according to the method recommended by Dole ( 1 ) against a potassium acid phthalate solution, which was prepared from salt obtained from the Sational Bureau of Standards according to directions supplied by the bureau. An additional check was made as recommended by Dole against sodium borate. The E” values were determined before each series of titrations, using Equation 1:
h concentrated synthetic sample of sea water having the following comoosition was oreoared from carefullv ourified salts: SaC1, 074920 hf; MgCln, d.0@6 M ; Sa2SO4, O.Q“3g8&I;. CaC12, 0.0138 M; KC1, 0.0123 M; haHCOa, 0.0026 M . Solutions for titration were prepared by dilution of this concentrated solution. Eight samples of natural and four samples of artificial sea water ranging in ionic strength from about 0.07 to 1.0, were titrated. The chlorinity range was from about 2.0 to 29 grams of halides per liter of sea water. The pH of the sea water varied between 7.8 and 8.2. During titrations the solution temperature was maintained constant Fithin l obetween 22” and 25” C. CALCULATIONS. The pH u’as calculated using Equation 1, the “glass electrode hydrogen-ion activity”, C H + ~ H +according , to Equation 2 pH = -log C E + ~ ~ H + (2) and the normality of the excess strong acid according to Equation 3 (ml. of excess HCl) (.V HC1) CH’ = (3) (ml. of sample ml. of HCl)
+
1 Present address, Department of Chemistry and Chemical Engineering, University of Idaho, Moscow, Idaho.*
Ezample.
767
A 100-ml. sample of sea water plus 25.93 nil. of
’
INDUSTRIAL AND ENGINEERING CHEMISTRY
768.
Vol. 18, No. 12
DETERMINATION OF ALKALINITY OF SEA WATER
Table I. Coefficient Data for Sample, CI = 18.83 Grams per Liter HCI, Ml. 23.93 24.13 24.33 24.53 24.93 25.43 25.93 26.92 29.94 34.91 39.90 44.80
PH
(CHi.fHf)(lOS) 4,266 5,248 6.397 7.465 9.638 12.62 15.49 21.38 38.11 63.83 88.92 112.5
(CH+)(1O5) 3.31 4.91 6.51 8.12 11.3 15.2 19.2 26.8 49.3 84.4 117.0 147.3
fH
Exactly 25.00 ml. of 0.01000 N hydrochloric acid are placed in a clean dry 135-ml. glass-stoppered bottle that has been aged for several months with hydrochloric acid having a pH of about 3.5. A 100-ml. sample of sea water is pipetted into the bottle, the sample is brought to the desired temperature, and the p H is measured. This measurement may be made aboard ship, but the treated sample may be safely retained and the pH measured more conveniently ashore. As indicated in Equation 7, t h e normality of the excess strong acid, CH,is calculated by dividing C H + ~ Hobtained + from the pH measurement of the solution by the appropriate coefficient, j H , for the water being analyzed.
(7)' Table II. Summary of
. Sample
C1, Grams
per Liter 2.09 2.84 7.03 9.60 12.31 15.38 15.86 16.97 18.83 10.56 23.44 28.93
{H+
pH Range 4.07-2.93 4.00-3.04 3.95-2.98 4,09-3,03 3,62-3.01 4.01-2.95 3.92-3,03 3.69-3.01 3,81-2,95 3.94-3.08 3.98-2.93 3.99-3.02
Data fH
'
0.844 0.798 0,764 0,757 0.748 0,746 0.763 0.766 0.777 0,763 0.777 0.768
Average Deviation 0,010 0.006 0.007 0.005 0.007 0,009 0.010 0.016 0.016 0.007 0.011 0.005
By substituting the C H + value in the following equation, developed by Thompson and Bonnar, the alkalinity may be calculated: 'Oo0
) (ml. of HC1) (-V HC1)1 -
Since all the above factors are k e d except the alkalinity and the CH+, Equation 8 becomes Alkalinity = 2.500 - (1250)
0.00999 N hydrochloric acid yielded a solution with a pH of 3.81. The equivalence point occurred a t 23.52 ml. of the acid. Therefore,
(v)
and it is possible to prepare tables or graphs which will relate pH or the C H + ~ Hand + alkalinity for a certain chlorinity range where no significant variation infH+ occurs. Further(4)
C H + ~ H= + 15.49 X 10-6 a t pH 3.81
(5)
I
'
The fH+ values were not constant throughout the pH range investigated; the variations in the values are given for one sample in Table I. The data for the twelve samples titrated are given in Table I1 and in Figure 1. The pH range over which the coefficients were sensibly constant is given for each sample. I n these measurements the limiting factor was pH. The reproducibility of the pH measurements was 10.005 pH unit, which amounts to a 1,2'%variation in the f ~ values. + The variation off^+ with pH is considered to be due to the decreasing ionization of the carbonic acid as the amount of excess hydrochloric acid increases. It is possible, making certain assumptions, to correct for the ionization of the carbonic acid. When this is done the f ~ +values a t the higher pH values are lowered and the results are in agreement with those a t lower values where the excess hydrogen ion is sufficient to repress the ionization of the carbonic acid. However, this must not be done for the present application of the fa+ results. I n Figure 1 the size of the circle indicates the extent of deviation of the individual measurements. From this figure the values off^+ in Table 111 have been taken.
Table 111.
Interpolated Values of f H t
Chlorinity, grams per liter Ionic strength Coe5cient
2 0,072 0.845
4 0.143 0,782
6 0.215 0.770
8 0.287 0.760
10 0.369 0.755
Chlorinity, grams per liter Ionic strength Coe5cient
12 0.430 0,752
14 0,502 0.752
16 0,574 0.754
18 0.646 0.764
0.717 0.768
20
I Y
C
,E ._ 0.9 : U
Y
'D
._ L ._
g 0.8
W
0.7
0
8
16
24
Initial Chlorinity per Liter
Figure 1 . Variation of Coefficient of Hydrogen Ion with Initial Chlorinity of Sample and Ionic Strength after Addition of 25 ml. of A c i d To test the proposed method three samples representing rt chlorinity of 18.83, 12.31, and 7.03 were treated with several different amounts of standard hydrochloric acid. The alkalinities were calculated and compared with the values obtained by titration; the results are shown in Table IV. These results indicate a probable error in the neighborhood of 1%, which is not an unreasonable or unusually high value, considering that the concentration of the solution is only about 0.002 molar with respect to the bicarbonate ion. The major factors that intluence the results of the determination have been investigated and the following conclusions
December, 1946
ANALYTICAL EDITION
obtained. If the pH varies by 10.01 unit a t 3.60 the alkalinity will vary by 10.5%. This same percentage variation results from a change of 1 3 ’ or a change of +2.5% in the value offH+. .-ln error in measuring the volume of the standard acid of 10.05 ml. and an error of 6 2 parts in 1000 in the normality of the standard acid produce a change of +0.257,, while an error of h O . 1 ml. in measuring the volume of the sample produces an error of only * 0 . 1 5 ~ oin the alkalinity values. At this laboratory the acid is measured with calibrated automatic pipets and the sample bottles are usually prepared before the ship is under way. The sample is measured using a measuring pipet described by Thompson and Anderson (4). The pH values usually encountered are around 3.6. Under abnormal conditions they may rise to 4.0 and higher, but this can be remedied by adding more acid, Tvhich is a desirable thing to do, as this brings the fa+ values into the region where they are more nearly constant. The method described for determining the alkalinity of sea ,water routinely should be applicable to a number of other systems, such as wash waters, effluent from manufacturing opera.tions, urines, blood, and other biological fluids. If the procedure is to be applied successfully to other systems, t h e test solutions must be fairly uniform as to ionic strength .and the concentration of the unknown material. Then by using a standard size sample and a fixed amount of reagent, the relationship between the constituent whose concentration is being determined and the pH, or any electrode e.m.f. t h a t is related t o the test substance, can be determined. Once the relationship has been established, the laborious titration process can be eliminated. Obviously, where only an occasional determination is to be made, no benefits are found. However, where large numbers of samples are examined in routine industrial or clinical work, a considerable saving of t,ime is effected.
Table
769
IV.
Sample
Summary
pH of Treated Sample
A
3.05 3 . 67 67 4.. . 1.. 3 4.37
B
3.01 3.39 3.91
C
2 98 3.47 3.65
of Calculated Alka’linity Values Alkalinity, Milliequivalents per Liter Calculated B y titration
Percentage Error
2.35 2.33 2.33 2 :. 3 2 Av. 2 . 3 3
2.35
1
Av.
1.74 1.72 1.70 1.72
1.71
0 5
av.
1.22 1.22 1.21 1.22
1.21
1
.
ACKNOWLEDGMENT
The authors are especially grateful to T. G. Thompson for his contribution to the preliminary phases of this investigation and to E. C. Tingafelter for his very valuable suggestions and criticisms. LITERATURE CITED
(1) Dole, Malcolm, “ T h e Glass Electrode”, New York. J o h n Wiley & Sons, 1941. (2) Mitchell, P. H., and RakeAtraw, N. W., Biol. Bidl., 65, 437-45 (1933). (3) Sverdrup, H. U..Johnson, M . W., and Fleming, R. H., “ T h e Oceans, Their Physics, Chemistry, and General Biology”, p. 193, New York, Prentice-Hall, 1942. (4) Thompson, T. G., and Anderson, D. H., J . Maririe Research, Sears Foundation, 3, 224-9 (1940). (5) Thompson, T. G., and Bonnar, R. U.,IND.k h ~ CHEX.. . ANAL. ED.,3, 393-5 (1931).
Accelerated Ozone Weathering Test for Rubber JAMES CRABTREE AND A. R. KEMP Hill, N. J.
Bell Telephone Laboratories, Inc., Murray
Light-energized oxidation and cracking b y atmospheric ozone are the agencies chiefly responsible for the deterioration of rubber outdoors. Since these processes are separate and distinct, i t is proposed to distinguish between them in the evaluation of rubber
for resistance to weathering. A n accelerated test for susceptibility to atmospheric ozone cracking is discussed. Apparatus for conducting the test and for measurement of ozone in minute concentration i s described in detail.
I
light-colored goods. Susceptibility can be readily gaged by measurement of the oxygen absorbed when the material is irradiated in air or oxygen under controlled conditions. The present obstacle to formulation of a standard test is the lack of a constant and reproducible light, source simulating sunlight. This article is principally concerned with the ozone factor. I n the case of rubber goods containing appreciable loadings of carbon black, such as insulated cable jackets and hose, by far the most, important cause of deterioration is the familiar cracking a t atressed locations by t,he ozone of the atmosphere, popularly and erroneously referred t o as “light cracking”. Many workers have attempted to use ozone-cracking in t h e laboratory as a measure of outdoor-cracking, only t o abandon the method because of failure to duplicate outdoor ratings. The probable cause of this failure has been found by the authors to be the use of too high concentrations of ozone. Small additions of wax, for example, which would confer substantial protection to cracking a t atmospheric concentrations of ozone, are quite Jyithout effect in most cases a t concentrations one hundred times greater. If ozone concentrations approximating that, of the atmos-
T H.\S been previously shown by the authors ( 1 ) that
tTyo
separate and distinct factors are responsiblc for the changes which rubber compounds undergo when exposed outdoors: a light-cuergized oxidation and attack by ozone which is normally prescmt in the atnmpliere in minute concentration. The former acts independently of stress in the rubber, the latter only when the material is stretched. Since sunlight, ozone, and temperature vary hourly, daily, seasonally, and with locality, the over-all result t o the exposed material will depend on t,he balance between the causative factors and n-ill be unique, since any given combination of such factors will not be duplicated. I t is therefore essential that the susceptihilities of a given compound t o damage by these two factors be dcterminrd separate1 The information thus furnished will then permit of an es ate of the durability of the compound under any givcn set of conditions. Damage by light-energized oxidation can be mitigated to any great cxtent only by physical protection from light, which can be partially achieved by incorporation of carbon black or ferric oxide. I t is likely to be the major factor in the case of
