MARCH, 1940
INDUSTRIAL AND ENGINEERING CHEMISTRY
FERTILIZERS AND MATERIALS
The full force of the war was felt directly by the traffic in potash fertilizers; the year's totals dropped sharply and accounted for most of the reduction in the fertilizer import trade as a whole. Table IV shows the extent of the reduction. Ammonium sulfate imports decreased from $3,088,000 in value in 1938 to $2,987,000 last year, and ammonium nitrate mixtures from $1,894,000 to $1,810,000. Calcium cyanamide imports, however, rose slightly, the 1939 value being $3,174,000 and that of 1938, $2,996,000. Sodium nitrate also registered a small gain; its value was $11,213,000 as against $10,732,000 in the previous year. OF POTASII FERTILIZERS (IN THOUSASDS) TABLEIV. IMPORTS
Chloride, crude Kainite Manure salts Sulfate Nitrate (saltpeter)
-1938Tom 199 54
Value
$5372
8
526 113
65 53
2193 1652
-1939Tons 84 19 2
48 58
Value
$2314 155
22 1624 1636
In the export trade ammonium sulfate and other nitrogenous chemical fertilizers were sold to a slightly smaller value,
423
$4,627,000 as against $4,735,000 in 1938; phosphate rock also declined from $6,638,000 to $5,233,000, while superphosphates registered $1,010,000 last year as against $945,000 in the preceding year. Potassic fertilizer materials rose to $4,447,000 from a 1938 figure of $2,600,000. EXPLOSIVES, FUSES, ETC.
This classification does not include explosives in the form of ammunition. The chief import is firecrackers, the value of which dropped from $613,000 to $377,000 last year. Exports are largely dynamite, the value of which in 1939 was $2,857,000, comparing with $2,186,000 in 1938. SOAPS AND TOILET PREPARATIONS
Perfume materials, $1,983,000 last year as against $1,453,000, made up the chief item in the import trade in 1939. Exports of this class of goods comprise a considerable number of soaps and toilet items, chief of which is dental creams. Trade in that commodity was valued a t $1,964,000 in 1939 as compared with $1,812,000 in 1938. Among the soaps, toilet or fancy soaps made up nearly half of the totai value, shipments increasing from $1,385,000 in 1938 tb $1,675,000 last year.
The pH of Sodium Dichromate Solutions
DILUTION EFFECTS
H. J. KAUFMA", W. B. LAUDER, AND R. K. KEPNER I S C E the manufacturers and consumers of dichromate are Mutual Chemical Company of America, Jersey City, N. J. interested in its pH, this investigation was conceived as an opening into what has been a more The pH of solutions of sodium dichromate varied considerably as or less untouched field. In studying the pH of colored solutions such as these solutions were diluted from approximately 4.5 to 0.05 mole dichromates, which are also strong per liter. The curves so obtained were but little affected by the oxidizing agents, glass electrode indistilled water used to make Lhe solutions, but were affected to ai struments are most satisfactory. greater extent by the container in which they were prepared and A survey of the literature reveals tested. Small changes in temperature had a negligible effect. no data taken directly for the purpose of determining dilution effects Concentrating rather than diluting the solutions had no effect in' of sodium dichromate. Much work the weaker ranges, but gave more acidic values in the stronger.! has been done, however, on titrating Additions of sodium sulfate and sodium chloride, two of the several chromic acid with sodium hydroxide. impurities normally present with sodium dichromate, were found, It is conceivable that the point a t to have but a small effect on the pH of the solutions. which 100 per cent sodium dichromate exists-in the solutions of these investigators might give some results comparable to the present work. However, recalculation of their data reveals that no conclusions can be so easily drawn. Furman (Id), Hughes (Is),and Margaillan (22) present their data in a manner difficult to recalculate to the basis of this paper, and Britton (3) worked a t concentrations too low for purposes may be derived the equations: of comparison. Consequently these data could not be properly compared with the present work and have not been included. applied the Debye-Huckel equaNews and Riemann (W) tion to chromate-dichromate mixtures in order to determine the second ionization constant of chromic acid. From the equilibria existing in such solutions,
S
INDUSTRIAL AND ENGIKEERING CHEMISTRY
424
VOL. 32, NO. 3
and collected in a silica beaker. This gave water of the required pH. A rapid determination of this p H was necessary to minimize the effect of alkali dissolving from the glass electrode. Many authors (4,7 , 11, 18) have considered the effect of the containing vessel on the p H of unbuffered solutions, and some of them state that glass is not a satisfactory material. This was shown by the fact that when water was stored for any length of time in glass bottles, the p H gradually rose. Furthermore distilled water from the tin condenser collected in a glass, rather than a silica, beaker had a pH higher than 5.8.
Preparation of Pure Sodium Dichromate and of the Solutions
MOLES
NarCrtOy PER L I T E R
FIGURE 1. SOLUTIONS OF c. P . Na2Crz07.2H20 WATERSAND CONTAINERS
IN
VARIOUS
Ordinary commercial sodium dichromate contains approximately 0.3 per cent of impurities, principally sodium chloride and sodium sulfate. This salt was recrystallized twice and yielded a product containing less than 0.003 per cent sodium chloride and 0.003 per cent sodium sulfate. The ordinary Mohr test was used for sodium chloride, and sodium sulfate was tested for gravimetrically after reduction of the dichromate by alcohol.
where parentheses represent activity, brackets represent concentration, and K' and K: are the apparent equilibrium constants. These equations will be later applied to the work of this paper. Extensive bibliographies may be found in such reference works as Clark (6),Britton ( W ) , and Dole (IO).
Instrument and Manipulation The Leeds & Northrup portable pH meter No. 7660 was used. This consisted of a potentiometer circuit with thermionic tube amplification. The glass electrode (4, 21) was composed of a thin glass bulb (Corning 015) filled with 0.1 N hydrochloric acid in which was suspended the quinhydrone assembly. The calomel electrode, containing saturated otassium chloride, had a groundglass sleeve as the liquid-liquif junction. The accuracy of the instrument as stated by the manufacturer was * 0.03 pH unit. This meter contained no temperature compensator, so that each solution had to be adjusted to a standard temperature, chosen as 25" C. Dichromate solutions within the range of this work showed an allowable temperature variation of 2.5" C. before a change in pH measurable by the instrument could be detected. Since no variation of more than 1' was ever experienced, a thermostatic control was unnecessary. A difficulty often experienced in measuring pH is the tendency of the electromotive force to drift until it reaches an equilibrium value. However, in working with the solutions described in this pa er, such drift was not found. 211 glass electrodes show a phenomenon known as asymmetry potential. This potential necessitates the application of a correction to any reading taken with the electrode. Furthermore, such a correction may vary from day to day. Consequently the electrode was standardized at least once each day with two or more buffer solutions of known pH to establish this correction.
pH of Water Since water is the solvent used in the solutions discussed here, its pH is a matter of some importance. This subject has been given serious consideration by numerous investigators ( I , 4, 6, 7 , 8, 12,13, 17-20). Various values have been reported for the p H of distilled water, but the accepted one is 5.8. The reason that such water does not attain a p H of 7.0 is probably that it reaches equilibrium with the carbon dioxide of the air. Since distilled water from the laboratory still had a pH of 6.5, i t was decided to prepare water of the p H designated in the literature-5.8. A condenser containing a tin condensing tube was set up, and the distilled water redistilled
\
290
2.85,
I
5 . 8 P H WATER '
1
BEAKER
1.5
,
l
~/'II
!
1
I
p
1 I
GLA53
IO
I
\d m
l
1 1
PO
1
?,5
3.0
I
35 4.0
M O L E S IdatCrz 01 PEE LITER
FIGURE 2. MAGNIFICATION OF THE CRITICAL REGIOX OF FIGURE 1 FOR 5.8 PH WATER
Commercial sodium dichromate generally contains sufficient sodium chromate to affect markedly its pH in water solution. Consequently to ensure complete removal of this material, the recrystallization was continued until the pH of samples taken from two successive crystallizations was the same. It was thereupon concluded that any material affecting the pH of the dichromate had been removed, and that the crystals contained 100 per cent Na2CrzOi.2H20(less the minute quantities of chloride and sulfate). The purpose of the in\-estigation was to see what relation existed between pH and concentration. Concentration was varied in two ways. First the saturated solution of pure sodium dichromate in the desired water was prepared (heating to dissolve the dichromate and then cooling to 25' C. and filtering out excess solid salt). This solution was then diluted with successive portions of water. Secondly, a weak solution was prepared which was successively concentrated by adding solid dichromate and maintaining the temperature a t 25 C.
MARCH. 1940
INDUSTRIAL AND ENGINEERING CHEMISTRY
425
The first method is called the "dilution" method, the second the 'Lconcentration" method. The strength of the solutions was checked by reducing the sodium dichromate with excess ferrous sulfate reagent and then back-titrating with potassium permanganate. Details of this procedure may be found in Scott (24).
Dilution Curves The curves in Figures 1 and 2 show the results of testing solutions by the dilution method in both Pyrex glass and silica beakers. Waters of 6.5 and 5.8 p H were used. Figure 2 is a magnification of the critical region of Figure 1 for 5.8 pH water only. The curves are similar in shape. They all tend to come together a t the lower concentrations-that is, below 1 mole per liter. The most striking fact, however, is that although all the curves vary above I mole per liter, greater variations appear between solutions tested in different containers than between solutions made from different p H waters. The effect of different waters was not negligible, as might be expected, but was quite small. The solutions tested in glass beakers show decided buffer effects in the region of 2 to 3 moles per liter, whereas the curves of solutions tested in silica beakers are sharper and more acid in this range. The explanation of this effect was not definitely determined though it is the conclusion of the authors, as well as of certain other investigators (7, 11, l a ) , that sodium oxide is dissolved from the glass. The alkali so dissolved would naturally neutralize some of the dichromate and produce a buffer mixture (a weak acid plus its sodium salt). Consequently for most accurate results i t would appear desirable to use silica rather than glass containers. The possibility that silica might be dissolved and make solutions in such containers more acid was checked by obtaining data in platinum. This gave the following deviations from the curve in silica. hloles NazCrzOr.2H20/Liter 1.13 1.68 2.68 3.64
PH -0.03
-0.04 +O. 03 +0.03
These values are within the limits of accuracy of the instrument and therefore indicate that the silica container has no effect on the dichromate solutions. I n preparing the solutions i t was necessary to stir them with glass rods and to filter them in glass funnels since silica apparatus of this kind was not available. However, on comparing pH values of solutions so prepared with those that had never come into contact with glass, no measurable differences could be found. This could be explained on the basis that high temperatures and long times of contact are the principal factors in the solution of glass. Both of these conditions were minimized above.
Concentration Curves Figure 3 is typical of the concentration runs. I n the region from 0.03-3.0 moles per liter, the concentration and dilution curves practically coincide, but from 3.0 up, the concentration curve is more acidic. The shift observed was common to all such runs made. It may be due to a change in the hydration of the ions since the solutions were prepared differently. Although this opinion has not been verified, the fact that the solutions are definitely dissimilar in character has been established. Other properties of these solutions were determined and discovered to differ. Vapor pressure, viscosity and conductivity all varied. These variations could be interpreted as indicating a change in hydration of the ions.
r a
io
I5
'2.0
25
"vICLf5 '!nZCrzO; PEE
F I G ~ R3.E PH
3 0 3 5 40 05 LITE.^
O F S O L r T I O S BY CONCEiTTR4'fION
One of the solutions prepared by the concentration method and one prepared by the dilution method were allowed to stand for a week; in that time neither showed any appreciable change in pH. This indicates that the solutions are stable. Since, however, the changes involved occurred only a t the higher less-important concentrations, it was felt, that a p proaching each curve from two directions was unnecessary. The shift, while appreciable, had little quantitative value for industrial purposes.
Discussion As has already been noted, the general shape of these curves is similar in all cases. Furthermore, they resemble dilution curves of other similar salts. Thus, disodium hydrogen phosphate (11) gives a similar curve except that, being a n alkaline salt, it has a maximum rather than a minimum pH. Also the reverse occurs a t a lower concentration. The reason that these curves bend back on themselves has been suggested in the literature as a change in the ionic strength of the solution (16) and possibly a change in the dissociation constant. Although Neuss and Riemann (23) worked with chromatedichromate mixtures of potassium rather than sodium salts, we shall nevertheless use their data to calculate values of the PH. Replacing [Cr04--] in Equation 4 by [H+] from Equation 2, replacing [H+] by (H+)/fx (fH = activity coefficient of the hydrogen ion), and expressing fa according to Debye and Huckel, (5)
pH = log l/(H+) y = [HCr04-] p = ionic strength of solutions a = a constant related to av. effective diameter of
where
(6)
ions
Neuss and Riemann give a value of 0.7 for 3.3 a 10'. The work of these authors was done a t lower concentrations than that of the present paper. Their strongest concentra-
INDUSTRIAL AND ENGINEERING CHEMISTRY
426
tion however ( p = 0.16) may be found on the curves of this papjer and so may be used for calculation. At p = 0.16, K ; := 9.51 X lo-’. By solving the following two equations, y may be found. If we neglect [H+]and [CrOd--l, JL
= 3c
- ‘/2y
(7)
From Equation 3, K’ = where C K‘
Y2 c - ‘/zY
concentration of sodium dichromate, moles/liter
0.0114
From this, y = 0.0228 and C = 0.0571. Then from Equation 5, pH = 3.91. From Figure 1, a t C = 0.057, p H = 3.89. This is good agreement. The extension of this calculation to any higher concentrations would be invalid, since the theory holds only at relatively low concentrations. The recorded pH values in concentrations over 1 mole per liter are probably not pH values as defined by Equation 6 but represent a combination of electrochemical effects. It is conceivable that several corrections might be applied, but since only a few of these corrections can be evaluated, it would be useless to apply any of them. Dole’s correction for water activity (9, IO) has been evaluated, but this correction alone proves insufficient to explain the shape of the curves. The value of measurements b a d e a t concentrations over 1mole per liter is that such measutements are reproducible and provide a sensitive means of checking purity, particularly between 2 to 3 moles per liter. Although they should not be interpreted as a direct measure of the hydrogen ion activity, they probably reflect changes in this property of the solutions.
Impurities The data for c. P. sodium dichromate is interesting but is of limited value in actual industrial use since teclinicel dichromate contains impurities. Of these only sodium chloride and sodium sulfate have been studied thoroughly. Other impurities which affect the pH are being investigated. The effect of these impurities could be studied best by preparing them in a pure condition and adding known quantities to a solution of sodium dichromate. I n all cases the solution was heated to dissolve the salts, cooled, filtered, and then tested for the impurity. SODIUMCHLORIDE. Ordinary analytical type salt was considered sufficiently pure for this purpose because it gave a solution with a p H of 5.84 when made up to 2 grams per 150 cc. of water. Three concentrations of salt were used-approximately 1.0, 0.5, and 0.05 per cent; i. e., the actual weight of sodium chloride present was 1 per cent of the weight of Ka&r?O,.2Hz0, etc. These tests wefe made by the dilution method with 5.8 pH water only. Tire differences in pH from that of pure sodium dichromate are measurable but so small and so scattered as to be of no practical importance in this investigation. SODIUM SULFATE. The preparation of a pure sodium sulfate involved considerable difficulty, Samples of c. P. sulfate f r o b various manufacturers gave great variations in pH; some were acidic and others basic. It was decided that a sulfate which crystallized from solutions of different pH values with the same pH could be considered a standard and reproducible material. Solutions of sulfate were therefore prepared and adjusted to various pH‘s. The pH’s of the resulting sulfates are as follows:
VOL. 32, NO. 3
Original Solution
Solutions of Recrystallized Sulfate
37
5.61 5.95 5.98 6.05
4.1
5.1 9.2
The above testa were made in solutions of 2 grams of sulfate in 150 cc. of water. All samples seemed to approach 5.95; therefore this was accepted as the pure sulfate. The p H of solutions of this salt was accordingly not greatly different from that of solutions of sodium chloride. Concentrations of approximately 0.1, 0.2, 0.5, and 1.0 per cent were used. These percentage values are calculated on the same basis as those in the sodium chloride determinations. The differences for the most part are no larger than those with sodium chloride, but they are practically all negative which indicates that they are of a definite nature. The average difference is 0.04 pH unit. MIXTURES. Sodium sulfate and sodium chloride never occur singly in dichromate but always as a mixture. Consequently a typical mixture, 0.2 per cent sodium chloride and 0.2 per cent sodium sulfate, was prepared and examined for pH. The differences were no greater than for sodium sulfate alone and therefore can be ignored.
Conclusions The p H of the distilled water used in making molality us. p H curves of sodium dichromate has only a small effect on the resulting curves, but the container in which the solutions are prepared has a larger effect. The shapes of all curves so obtained are similar in that the observed pH values of all pass through a minimum a t a concentration of about 2.5 moles per liter. Small temperature variations have a negligible effect on the pH value. Curves formed by concentrating a weak solution and those formed by diluting a strong one almost coincide except at the highest molalities. Of the impurities present in technical sodium dichromate, only sodium chloride and sodium sulfate were studied. These, however, have a negligible effect on the pH.
Literature Cited Beans and Oakes, J . Am. Chem. SOC.,42, 2116 (1920). Britton, H. T . S., “Hydrogen Ions”, 2nd ed., London, Chapmau and Hall, 1932. Britton, H. T S., J . Chem. SOC.,125, 1572 (1924). Burton, J. 0.. Matheson, H., and Acree, S. F., Bur. Standurda J. Research, 12, 67 (1934). Byck, H. T., Science, 75. 224 (1932). Clark, W. M., “Determination of Hydrogen Ions”, 3rd ed.. Baltimore, Williams and Wilkins Co., 1929. Cliquet-Pleyel, R., Documentation sei., 5 , 65-70, 161-70 (1935). Cocking, T. T., Ind. Chemist, 12, 299-301 (1936). Dole, Malcolm, J . Am. Chem. SOC.,54, 2120 (1932). Dole, Malcolm, “Measuring pH with the Glass Electrode”, Maywood, Ill., Coleman Electric Co., 1937. Ellis, S. B., and Kiehl, S. V., J . Am. Chem. SOC.,57, 2139 (1935). Ibid.. 57, 2145 (1935). Fulmer, E. I., Iowa State Coll. J . Sci., 1, 37 (1926). Furman, N. H., J . Am. Chem. SOC.,44, 2685 (1922). Hughes, W. S., Ibid., 44, 2860 (1922). Kiehle and Loucks, Trans. EkctTochem. SOC.,67, 81 (1934). Kling, A.. and Lassieur, A., Compt. rend., 181, 1062 (1925). Ibid., 189, 637 (1927). Kling, A., and Lassieur, A,, Documentation sci., 4, 225-9 (1935). Lassieur, A., Ibid., 5, 11-15 (1936). MacInnes, D. A,, and Dole, Malcolm, J . Am. Chem. SOC.,52, 29-36 (1930). Margaillan, H. L., Compt. rend., 157, 994 (1913). Neuss, J. D., and Riemann, W., J . Am. Chem. SOC.,56, 2238 (1934). Scott, W. W., “Standard Methods of Chemical Analysis”, 4th ed., pp. 162, 1368f, New York. D. Van Nostrand Co., 1929. Thompson, M. S., Bur. StandaTds J . Research, 9, 833 (1932).
