Determination of l o w Alkalinity or Acidity in Water T. E. LARSON and LAUREL HENLEY, ///;nois State Water In determining mineral constituents of rain water, a precise determination of extremely low concentrations of alkalinity and acidity was essential. The procedure developed is based on the principle that increments of added acidity (after neutralization) increase the hydrogen ion concentration as a linear function. Extrapolation to 1 X 10-7 mole (H+) per liter provides a precise equivalent end point. The procedure is sensitive and accurate to 0.05 p.p.m. (as calcium carbonate) when titrating a 200-ml. sample of -1 to f l p.p.m. alkalinity with 0.02Ri sulfuric acid. I t might also be used for analysis of deionized water, steam condensate, or air pollution samples.
T
ITRATIONS of strong and weak alkalies with strong acid have long been used in quantitative analysis for standardization with primary standards and for the routine determination of alkalinities in natural waters (1). The use of methyl orange and mixed indicators has proved satisfactory for titration of alkalinity concentrations sufficiently high to ensure a reasonably definite p H a t neutralization. In the analysis of rain water samples in connection with the Cloud Physics Project, sponsored by the U. S.Signal Corps, the concentrations of alkalinity or acidity were found to be low but often a significant portion of the total mineral content. These samples may contain more than 10 times as much carbon dioxide as alkalinity. Therefore, the end point p H is not only variable by definition but unstable owing to loss of carbon dioxide during titration. The difficulty is similar to the problem of the titration to the phenolphthalein end point for free carbon dioxide in waters of very high alkalinity. Attempts were made to add predetermined concentrations of alkalinity to the samples, but the absolute error introduced by the measurement of added alkalinity and lack of precision in detecting the end point prohibited results of the required precision. It was then reasoned that if a sample is aerated with carbon dioxide-free gas, the percentage of free carbon dioxide liberated by successive additions of acid would increase as the p H decreased. This was followed by the consideration that following neutralization, successive additions of acid should increase the hydrogen ion concentration (or activity) as a straight line function. Then by extrapolation to zero or to 1 x 10-7 mole hydrogen ion concentration per liter, the exact neutralization point could be established. Calculations revealed that extrapolation to zero rather than 1 X lo-' mole hydrogen ion per liter on a 200-ml. sample would introduce a positive error of only 0.005 p.p.m. alkalinity (Figure l), well within the desired accuracy and experienced precision. Calculat+ns further revealed that aeration was not necessary if the titrations were continued to p H values sufficiently low to suppress the ionization of carbonic or other weak acids to an insignificant proportion. For carbonic acid, ionization is suppressed to 10% a t about p H 5.36 and 1% at 4.36. Silicic acid is ionized 0.1% a t p H 7 and therefore exists essentially as unionized silicic acid a t this p H and lower, and causes no interference. For acetic acid the ionization is suppressed only to 10% a t pH 3.55 and 1% a t 2.55. The absolute error in hydrogen ion activity, as calculated from a possible error of 0.01 p H unit, increases logarithmically or 10fold per unit decrease in pH. Therefore, greater accuracy and sensitivity are obtained by intermittent aeration during the titration of bicarbonate alkalinity in order tha.t the extrapolation can be made from maximum values of pH.
Survey, Urbana,
111.
Theoretically, the slope of the extrapolation line should correspond to the activity coefficient for the hydrogen ion. This was seldom found to be true, because of the impossibility of an exact standardization of the p H meter. Lack of such rigid accurary in the standardization of the p H meter, however, did not influence the reliability of the determination, since the extrapolation TTas made to zero. The apparatus used consists of a Beckman Model G glass electrode pH meter, a magnetic stirrer, air diffuser, 5-ml. microburet, graduated in 0.01 ml. with a platinum capillary tip, and a water bath maintained a t 25" It 1' C. The sample holder in some cases was a borosilicate glass beaker and in others a platinum dish. No significant loss in accuracy was noted when beakers were used a t this temperature a t p H less than 7 . Stirring and air diffusion were continued for 2 minutes after the addition of each increment of acid and stopped during the period of pH determination. Typical results on rain water samples are shown in Figure 2. The number for each extrapolation refers to diflerent samples. Numerically, alkalinity is defined as the equivalent concentration of titratable base and is determined by titration with B standard solution of a strong acid to certain equivalence points. In this procedure the equivalence point is taken to be a t PH 7. The alkalinity is therefore equal to the equivalent concentration of titratable stronger base relative to a weaker acid, when using a standard solution of a strong acid. Therefore, ammonium acetate titrates fairly closely to an alkalinity equivalent of the ammonium ion concentration. Ammonium chloride has no alkalinity. The acidity of a water is caused by carbon dioxide, mineral acids, and salts of strong acids and weak bases. Silicic, carbonic, and acetic acids are not as strong as mineral acids and therefore show no mineral acidity by this procedure. Acetic acid, however, can be titrated to a sufficiently low pH t o expel carbon dioxide and back-titrated with excess 0.02N sodium hydroxide until complete ionization of acetic acid (99% a t p H 7 ) is approached. Subsequent increments of sodium hydroxide then produce equivalent increases in OH- activity. A straight-line extrapolation or to 0 mole (OH-) per liter then indicates the concentrato tion of weak acid. Care is essential to avoid reabsorption of carbon dioxide. Also, if ammonia is present it should be recognized that it becomes increasingly volatile with increase in pH, just as carbon dioxide is volatile a t low pH. pgyn. Alkalinity
0
IO
I)I
.03
(08
Ca COS) .04 .05
0
-
0
e
6 0
I
-4
2
0 ml. 0.02N H,SO, Figure 1. Calculated error introduced by mole extrapolation to zero instead of 1 X (H+) per liter for equivalence point 85 1
852
ANALYTICAL CHEMISTRY
Acidity IO 00 500
IO
20
Rp.m. AlkaWty (as Ca CO,) 30 4 0 50 6 0 7 0
bonk acid). The increasing solubility of glass a t higher p H values limits the reliability of this application.
43
80
CONCLUSIONS
400
45
- 300
m 0
‘0
z
X
f-m o
47
100
50
2
60 70 18
0
Figure 2.
2
4
6
8 IO 12 d. 0.02 N \SO,
14
16
Extrapolations for typical rain water samples
Numbers (38 to 48) refer to laboratory sample numbers
Further consideration will show that both weak and strong acid concentrations in the same sample can be determined by titration first with 0.02N sulfuric acid and extrapolating to zero (H+)for mineral acidity. This may then be followed by backtitrating with excess increments of 0.02N sodium hydroxide and extrapolating to zero (OH-) per liter. The difference between the end points represents the weak acid acidity (exclusive of car-
The procedure described for the determination of alkalinity has been found to be sensitive and accurate to fO.05 p.p.m. alkalinit!: or mineral acidity when present in extremely low concentrations. A discussion has been provided on the titration of strong baseweak acid mixtures and the estimation of respective concentrations of strong acid and weak acid in acidic samples. This procedure was devised for analysis of rain water. I t is equally applicable for analysis of steam condensate and impurities in deionized water. Alkaline steam condensate samples should be collected with a known quantity of acid in order that ammonia is not volatilized before and during the titration. The procedure may also be adapted to analysis of .air pollution samples. It is limited to low concentrations and by the care and technique of the analyst. It can be made more sensitive and accurate by use of a more sensitive pH meter and a more accurate measurement of titrant volume. LITERATURE CITED
(1) Am. Public Health Assoc., New York, “Standard Methods for the Examination of Water and Sewage,” 9th ed., 1946. RECEIVED for review September 1 4 , 1954.
Accepted December 29. 1954.
Correction Factors for Comparing Activities of Different Carbon-1 4labeled Compounds Assayed in Flow Proportional Counter M. L. KARNOVSKY, 1. M. FOSTER, L. I. GIDEZ, D. D. HAGERMAN, C. V. ROBINSON, A. K. SOLOMON, and C. A. VILLEE, Biophysical Laboratory and Department o f Biological Chemistry, Harvard M e d i c a l School, Boston, Mass.
A comparison has been made of the activities of several organic compounds counted as such or as barium carbonate obtained after combustion. In the windowless gas-flow counter, under the conditions specified, the relevant correction factors are much smaller than those reported in the literature for end-window Geiger counters.
I
K experiments with carbon-14 it is frequently necessary to
compare the activities of different substances. For example, a comparison might be desired of the activity of a labeled substrate with the activity of carbon-14 dioxide derived from it metabolically or by chemical degradation. Correction factors have been reported to be necessary in making such conversions of the activities of organic substances (counted as such) to equivalent activities determined by counting the carbon-14 as barium carbonate. The following factors have, for example, been reported in experiments in which end-BTindow Geiger Muller counters were used: Glucosephenylosazone Pyruvic-2,4-dinitrophenylhydrazone Wax
1.34 1.26 1.29
(5) (1) (8)
Over the past several years, the authors have made similar measurements on a windowless gas-flow proportional counter. With this counting technique and using thin samples mounted on stainless steel planchets, correction factors are found under routine conditions of counting in this laboratory to be significantly lower than correction factors previously reported. APPARATUS
The counter, used for these studies and described previously
( J ) , is of the flow proportional type and uses a gas mixture of
argon with 5% carbon dioxide. The dimensions are shown in Figure 1. The counter couples directly (without a cable) into the scaler-amplifier unit (Model 162, Nuclear Instrument and Chemical Corp., Chicago, Ill.) which is operated a t about 1650 volts with a 5-mv. input sensitivity. The efficiency for thin carbon-14 samples is about 40%. THEORETICAL
If a series of planchets of various thicknesses is made up from a single batch of material and counted, a curve obtained for counts us. sample thickness is similar to that illustrated in Figure 2. This curve is characteristic of the sample material and can be used t o make relative corrections between planchets of a given compound by correcting all counts to the same weight of sample. To determine the ratio of activities between an organic compound and barium carbonate, the self-absorption curves must be known for each and the relative scales of the curves-Le., the correction factor-must be determined. One wav to do this is to burn thk organic compound to obtain barium carbonate with the same activity per carbon. Self-absorflion curves may then be obtainedfor ANOJE the organic compound T and for the barium carCAThODE bonate derived from it. E These curves may be related to each other in scale by adjusting the counts per minute in E SAMPLE each case for the weight fraction of carbon in :1 OLANCHET Tf each compound. Ratios of counts obi 1-9 mm-1 tained from barium car1-14 3 mm.bonate and those from -1 19 mm.the organic compound may be calculated from Figure 1. Windowless proporthese curves for various tional flow counter ~
L
i ,
‘I
