Volumetric Determination of Aluminum Using Sodium Citrate ALBERT C. TITUS
AND
MELVIN C. CANNON, University of Utah, Salt Lake City, Utah
I
T HAS long been recognized that the aluminum ion will
react in the cold with excess of alkali citrate to liberate two hydrogen ions for each aluminum ion. I n the presence of the citrate the liberated acid is conveniently titratable. With an alkali tartrate the aluminum ion in the cold liberates three hydrogen ions, the number ultimately formed on complete hydrolysis of the aluminum in water. It is difficult to titrate the acid liberated during normal hydrolysis, even though the added base tends to cause completion of the hydrolysis, but very easy to titrate the acid liberated in the presence of the tartrate. Thus White (4)many years ago outlined a method for the volumetric determination of aluminum in its sulfate by the utilization of both these organic salts, and successfully applied it to acidic and basic sulfates. He found that barium hydroxide was better than sodium hydroxide as a standard solution for titrating the liberated acid. This base was standardized against sulfuric acid which contained approximately enough dissolved precipitated aluminum hydroxide to duplicate the conditions in samples being analyzed, thus eliminating a slight overtitration otherwise obtained in the “citrate step” outlined below. Starting with 3 grams of aluminum sulfate dissolved in 100 cc. of water, two 25-cc. samples were removed for titration. To the first was added a 50-cc. portion of neutral 10 per cent potassium sodium tartrate. It was then titrated with the barium hydroxide t o the phenol hthalein end point, thus neutralizing the acid equivalent to &ree times the number of moles of aluminum ion as well as any free acid. If free base was present, the volume used to titrate the liberated acid was diminished by an amount of base
equivalent to the free base. The second 25-cc. portion was evaporated to dryness, dissolved in 50 cc. of neutral 10 per cent sodium citrate solution, and after being allowed t o stand 10 minutes was titrated cold with the base t o the phenolphthalein end point. The amount of base used here was equivalent t o the acid formed in the liberation of two hydrogen ions for each ion of aluminum, plus the free acid or minus the acid equivalent t o the free base. Simple calculation enabled the amount of aluminum t o be determined, along with the amount of free acid or free base present. Thus the aluminum itself was determined by multiplying the volume difference between the two titrations by three and then calculating the amount present by appropfiate use of the molarity of the base. The alkaline earth sulfate did not immediately precipitate under the conditions of the experiment.
A study of White’s work (4) indicated that the mole ratio of citrate to aluminum ion was somewhere between 3.2 and 8.6, depending on the states of hydration of his compounds, which were not definitely given. Evaporating the second portion to dryness gave results which otherwise could be attained only after the solution containing these substances had stood for many hours. Recently Pavlinova (S) showed that thymolphthalein was a better indicator than phenolphthalein, for the titration of the acid liberated during the reaction of citrate with aluminum ion, since phenolphthalein changed color too quickly and thus gave a negative error. Using potassium aluminum sulfate of known molarity, she &st titrated sodium citrate with sodium hydroxide to the thymolphthalein end point, then added the alum solution, and again titrated in the cold until the indicator changed color. As long as the mole ratio for citrate to aluminum ion was not much less than 1.2, and up to at least a ratio of 4.1, two hydrogen ions were liberated for each one of aluminum.
She also found that neutral solutions resulting from the titrations became basic upon boiling and then cooling, which showed that less acid would be liberated were this heat treatment applied to the original citrate-alum solution before titration. On the other hand with a very small ratio the neutral solutions became somewhat acidic on heating. No attempt was made to plot the liberated acid accurately against the mole ratio for citrate to aluminum in experiments in which heat treatment occurred, although it was obvious to the investigator that the curve would have an interesting course. Different temperatures also had different effects. She found it necessary to add phenolphthalein to the thymolphthalein in the experiments in which heating took place, since heat affected the latter. In the authors’ preliminary studies with the citrate-aluminum reaction equilibrium was slow to be attained in the cold, but heating caused i t to be realized quickly. For this reason they made no attempt to evolve an aluminum method in which heating did not have a place.
Variation in the Mole Ratio of Citrate to Aluminum In deciding the amount of sodium citrate to be added, the effects of varying amounts upon a constant amount of aluminum were studied. The citrate-aluminum solution was f i s t heated and then cooled before the liberated acid was titrated with sodium hydroxide. In these experiments the general method given below was followed except as to the amount of citrate used, starting where portion I1 is treated. The results indicated that the mole ratio of citrate to aluminum should be kept within certain limits in the general method, since one could no longer rely on the liberation of just two moles of hydrogen ion for each mole of aluminum ion. The heating had destroyed this simple relationship, although in the authors’ opinion the gain in speed in attainment of equilibrium compensated for this. Using 3.35 millimoles of aluminum as the sulfate, the amount of citrate was increased to more than 90 millimoles in experiments in which the liberated acid was titrated with 0.496 molar sodium hydroxide. The mole ratio of citrate to aluminum thus varied from much less than 1up to more than 26. Complete hydrolysis of the aluminum would give 10.05 millimoles of hydrogen ion, which would be neutralizable by 20.26 ml. of the standard base. On the other hand the liberation of two hydrogen ions for each one of aluminum would result in the use of only 13.51ml. of base. With the small ratio of citrate to aluminum consequent upon the use of the smallest quantities of added citrate, the liberated acid approached the more than 20 ml. expected from complete hydrolysis, but as the mole ratio approached 1.2, the lowest value in the work of Pavlinova (S),there was a steep drop to a little over 50 per cent of that expected from complete hydrolysis. This value was less than the 66.7 per cent expected by the “cold method” of Pavlinova (S), or the method of White (4). Qualitatively Pavlinova had predicted such results but had not shown the exact influence of variation in added citrate. The curve representing volumes of sodium hydroxide used, rn ordinates, plotted against mole ratios of citrate to aluminum ion, as abscissas, kept on descending until a value of 137
138
INDUSTRIAL AND ENGINEERING CHEMISTRY
about 45 per cent of the acid expected from complete hydrolysis in the absence of citrate was obtained. At this stage the mole ratio was 1.7 and about 9.3 ml. of the base were needed. The curve then rose and from a ratio of about 3 to about 11 it was a straight line with a steep slope. From a ratio of 15 to 26 only two or three points were obtained, but the curve was almost horizontal in this region and appeared to be straight again. Here the liberated acid amounted to close to two hydrogen ions for each aluminum ion and thus for the first time the results remained, over a long range, similar to those of the earlier workers with their different experimental conditions. The authors’ “hot method” gave similar results when they used a ratio of 15 to 26, while Pavlinova (3) had used 1.2 to 4.1 with her cold method, and White (4) had utilized a ratio somewhere between 3.2 and 8.6, using a still different variation. White used the liberation of acid by the citrate as only part of his general method, but nevertheless relied on liberation of exactly two hydrogen ions for each aluminum ion. The authors deemed it unwise to attempt to establish a method that required such a large amount of citrate as needed with a ratio of 15 or more for citrate to aluminum ion. It was decided that further studies could profitably be made in the region lying between ratios 3 and 11, where the rapidly ascending curve formed a straight line. Here the effect of the citrate on a given amount of aluminum was variable but was proportional to the amount of citrate present and so to the mole ratio of citrate to aluminum. By using a constant amount of citrate and varying the aluminum, keeping the mole ratio within the above limits, and titrating the liberated acid another straight-line curve could be obtained when the base used in the titrations was plotted, as ordinate, against the amount of aluminum present.
Establishment of Aluminum Method
It was necessary to select an appropriate amount of sodium citrate for addition to the aluminum in all cases. Since less than two ions of hydrogen were liberated for each one of aluminum, an equation was evolved for calculating the amount of aluminum in the sulfate from the amount of base used in titrating the acid liberated. A slightly different equation was necessary with potassium aluminum sulfate, because of the influence of the potassium ion. Confirmation of this effect of potassium was qualitatively obtained by tests in which potassium chloride was added in varying amounts to aluminum sulfate samples, with consequent increase in titratable acid with constant amount of citrate and aluminum. Sodium chloride did not have this important effect on the acidity, The two equations were obtained by study of the curves made by plotting the base used against the aluminum present in the aluminum sulfate and the potassium aluminum sulfate. Reagents and Apparatus Ordinary distilled water was used. Reagent solutions were properly protected from carbon dioxide after being made up. Since the presence of carbonate in the sodium hydroxide did not matter in this work, it was desired only to avoid further absorption in the standardized solutions. Sodium hydroxide, 0.5 molar, was always standardized against two different samples of standard hydrochloric acid made up from separate sources of the constant-boilingacid. Sodium citrate, 0.610 molar, was made from freshly opened sodium citrate dihydrate which contained 11.8 per cent of water, according t o a private communication from the manufacturers. This U. S. P. product was later found t o be definitely acidic to the authors’ indicator when 4.7 grams were dissolved in a total volume of 100 ml. Under these conditions there was not over the equivalent of 0.06 ml. of 0.5 molar sodium hydroxide of ti-
VOL. 11, NO. 3
tratable acid. Titration was made in the cold, without previous heating. In making u sodium citrate solutions any sample not departing from the $ s. P. specifications of 10 t o 13 per cent of water is satisfactory, since an error of 2 per cent in the amount of citrate used is inap reciable. The theoretical water content is 12.25 per cent for t i e dihydrate. If the sample departs from the acidity specifications outlined above, correction must be made by adding sodium hydroxide or sulfuric acid. The authors added 1 ml. of carbon tetrachloride t o each liter of solution for preservative. Aluminum sulfate, AL(S04)&3H20, was dissolved, filtered, and found gravimetrically to contain 0.1815 mole per liter. This solution was used during tests in which the correction for free acid was made and thus added acid was tested for in various samples and satisfactorily recovered. Any free acid present in the sample itself would thus have been detected. Potassium aluminum sulfate, c. P., was found gravimetrically t o be 0.1776 molar. Potassium fluoride. Reagent quality material was used. Sulfuric acid was 0.5 N and was standardized against the sodium hydroxide. It was used, with the potassium fluoride, during tests on the procedure accorded portion I in the general method. Indicator solution. After exhaustive tests on various indicators present during electrometric titrations 0.075 per cent thymol hthalein and 0.025 per cent phenolphthalein by weight, in alcokol, were selected as indicator. Fifteen drops were used in the 100-ml. total volume and the “colorless to rose” indication was used.
Determination of Aluminum in Absence of Iron Sodium hydroxide containing carbonate may be used except when solutions are to be heated, Thus in establishing this method and applying it to neutral aluminum solutions, i t is not necessary to free the base from carbonate before standardization. Decidedly acidic solutions require adding a large corrective volume of base to portion I1 after making the determination in portion I. If this base contains carbonate, error is involved in heating portion I1 containing the added base. On the other hand adding the base after cooling the solution means heating the solution in the presence of the original free acid. With large amounts of free acid this is not permissible, as it would cause a shift in the end point of the ensuing titration. Thus with large amounts of free acid, and perhaps even with very small amounts, the corrective base in portion I1 must be added before heating and for this reason carbonate-free base is essential. Carbonate present in the sodium hydroxide will not cause error in the absence of free acid, and may cause no appreciable error with only fairly small amounts, provided the base is added after and not before heating. GENERAL METHOD.Start with an iron-free sample containing between 80 and 430 mg. of aluminum (150 and 810 mg. of the oxide), and make up to 50 ml. in a volumetric flask. Remove 25 ml. with a ipet and use as portion I. Quantitatively transfer the remaingr to another vessel as portion 11. Portion I. To find the correction for free acid add 6 grams of potassium fluoride, titrate with standard 0.5 N carbonate-free sodium hydroxide to the phenolphthalein end point, and discard the sample. [The treatment accorded portion I was based on the slightly modified method of Craig (1, 2). Tests made with the aluminum sulfate solution containing added sulfuric acid were satisfactory, in that the added acid was recovered. None of the solutions included in Tables I, 11, and 111 were treated in this way, as simple aluminum salts were used in establishing and testing the method and the alum solution was made from c. P. material while the aluminum sulfate solution was indirectly tested as directed above (see also “Reagents and Apparatus”). The use of phenolphthalein in portion I, where a few drops of 0.1 per cent indicator were used, is defensible, since here one is not titrating acid liberated from the reaction forming the citrate-aluminum complex. The mixed indicator should be as good as or better than the phenolphthalein here, unless it is harder t o get a sharp end point. I Portion 11. Add 25 ml. of 0.610 molar sodium citrate (see “Reagents and Apparatus,” discussion concerning making up and
MARCH 15, 1939
ANALYTICAL EDITION
testing citrate). Dilute to nearly 100 ml., add the volume of sodium hydroxide used for portion I, bring the volume t o 100 ml., and heat at the boiling oint for 5 minutes. Cool, and titrate to a light color with the gase, using the thymolphthalein-phenolphthalein indicator previously described. (The base need not be carbonate-free where there is no free acid correction.) METHODSOF CALCULATION FOR ALUMINUM PRESENT IN PORTION 11. For pure aluminum salts (potassium absent), Millimoles of A1
=
millimoles of NaOH X 0.697 - 0.341 (1)
For potassium aluminum sulfate, Millimoles of AI
=
millimoles of NaOH X 0.674
139
ADEQUACY OF EQUATIONS. In view of the fact that the data given are not shown in the form of the curves actually used, it seemed well to present in tabular form figures to show just how well the equations fit the data from which they were derived. Thus the last columns of Tables I and I1 present a comparison of the amounts of aluminum actually present with the amounts calculated by application of the equations. The percentage deviation is a measure of the adequacy of the equations, which are those fitting straight lines. This deviation, presented in the last column, indicates that over a long range the equations are applicable but that with smaller values for aluminum oxide they do not apply.
- 0.317 (2)
Discussion
In each case millimoles of aluminum may be converted to milligrams of aluminum oxide by multiplying by 50.97.
A sufficient test of the method of analysis, including the validity of the equations, is given above. The further test on aluminum sulfate merely confirms the finding for that mateTABLEI. DATAFOR EQUATION USEDFOR ALUMINUMSULFATE rial previously presented. SOLUTIONS This is essentially a repetition of the type of work used to (Equation 1 is derived from these data) obtain the original equation for aluminum sulfate. EquaNaOH = 0.494 molar. AIz(S04)a 0.1815 molar tion l is used in calculating the aluminum in solutions to Deviation which known amounts of aluminum have been added. The of Calod. Corrected Volume" A1 from calculated values are then compared with the known values as Ala(S04): NaOH A1 A1108 Calcd. KnownAl NaOH was done in Tables I and 11. In Table I11 the last column MI. MZ. Moles X 108 MQ. Moles X 100 % presents the percentage error of the calculated values ob2.88 2.01 1.423 0.730 37.2 0.651 -10.8 (*0.02) (AO.01) tained from the results of titrations when compared to the 6.29 5.01 known aluminum content. Again the authors began work 3.107 1.819 9 2 . 7 (10.03) 1.825 + 0.3 (10.04) 11.58 10.03 where portion I1 is treated, since the same aluminum sulfate (10.08) ( ~ 0 . 0 3 ) 5 721 3.641 185.6 3.647 + 0.2 16.87 15.02 solution was used as in the previous work. (10.07) (10.07) 8 . 3 3 4 5.452 2 7 7 . 9 5.468 + 0.3
-
22.26 (t0.02) a
20.05 .(*0.02)
10.997
7.278 371.0
7.324
+ 0.6
Four runs made in each case.
TABLE111. TESTINGMETHODOF ANALYSISWITH ALUMINUM SULFATE
ESTABLISHING CURVESFROM WHICHEQUATIONS WEREDESimple aluminum salts were used in establishing the curves from which Equations 1 and 2 were derived. The authors began where portion I1 is treated in the general method, since corrections for free acid were superfluous. RIVED.
TABLE11. DATAFOR EQUATION USEDFOR POTASSIUM ALUMINUMSULFATE SOLUTIONS (Equation 2 is derived from these data) NaOH = 0.494 molar. KAl(SO4)z = 0.1776 molar Corrected Volumea NaOH KAl(SO4)z MI. MI. 2.01 (10.01) 5.01 (10.03) 10.03 (10.03) 15.02 (10.07) 20.05 (10.02) 25.07 ( *O. 07)b 30.05
Deviation of Calcd. from KnownAl
AlnOa MQ.
A1 Calcd. Moles X 108
0.357
18.20
0.226
-37
1.719 0,890
45.36
0.842
-
1 . 7 8 1 90.78
1 764
4.416 2.668
136.0
2.659
3.561
181.5
3.548
NaOH
A1
Moles X 100 0.805
3.088
5.735
~
%
5.4
- 1.0 - 0.3 - 0.4
7 . 0 7 9 4.452
226.9
4.454
0.0
(AO.00)
8 . 4 3 3 5.337
272.0
5.367
f 0.6
(10.00)
9.776 6.239 318.0
6.272
+ 0.5
35.13
Four runs made. b Discarded a value 14.53 corresponding t o 25.07 in the next column, so only 3 runs. 0
The first five columns of Tables I and I1 contain the data used in plotting the curves for aluminum sulfate and potassium aluminum sulfate. The equations derived from the curves may be used in calculating the aluminum content of unknown solutions which are free from iron or from other interfering ions which liberate acid from citrate. In each table the numbers within parentheses refer to the maximum deviations of the volumes in individual runs from the average volumes given.
Run No. 1
2 3 4 5 6 7 8 9
10
11 12 13 14 15 16 17 18 19 20
(Equation 1 is used) NaOH = 0.494 molar. Alz(S0i)a = 0.1818 molar Corrected Volumes A1 A1 NaOH AIz(S04)r NaOH Found Contained MI. MI. Moles X 100 Males X 100 2.88 2.89 2.86 2.89 6.39 6.44 6.29 6.25 11.63 11.59 11.58 11.54 16.94 17.01 16.84 16.91 21.18 22.24 23.21 24.09
2.01 2.01 2.01 2.00 5.04 5.01 5.01 4.98 10,05 10,oo 10.04 10.03 15.06 14.97 15.09 14.96 19.07 20.07
21.06
22.06
1.423 1.428 1.413 1.428 3.157 3.181 3.107 3.088 5.745 5,725 5.721 5,701 8.368 8.403 8.319 8.354 10.463 10.987 11.466 11.900
0.651 0.654 0.644 0.654 1.859 1,876 1,825 1.811 3.663 3,649 3.647 3.633 5.491 5,516 5.457 5,482 6,952 7.317 7.651 7.953
0.730 0.730 0.730 0,726 1,830 1,819 1.819 1,808 3,648 3.630 3,644 3.641 5.467 5.434 5.478 5.430 6.922 7.285 6.645 8.008
Error
% -10.8 -10.4 -11.8 9.9 4- 1 . 6 f 3.1 0.3 0.2 0.4 0.5 0.1 0.2 0.4 1.5 0.4 1.0 i- 0 . 4 0.4 0.1 0.7
-
+++ ++ +++ ++ -
From the plotted data of all three tables the limits of applicability of the method were derived. The general method may be considered precise to 1 per cent over the region advocated for its use, which is between 150 and 810 mg. of aluminum oxide in the whole sample, or half of these values in the part used as portion 11. As variations within the prescribed limits of the acidity of the citrate used may account for 1.1 mg. possible error in evaluating the amount of oxide in an unknown, even the smaller samples used in portion I1 will be essentially within this precision limit.
Conclusion While sodium citrate is an essential reagent in an excellent volumetric aluminum method developed long ago, and is also used in a more recent cold method, the authors were not satisfied that equilibrium was attained quickly enough without heating. Their method involves a correction for any free
140
VOL. 11, NO. 3
INDUSTRIAL AND ENGINEERING CHEMISTRY
acid, and the cold titration of the acid liberated from a previously heated mixture of sodium citrate solution with the unknown aluminum solution. The ca,lculation of the ahminum content is based on substitution in an appropriate equation of the amount of sodium hydroxide used in titrating the liberated acid. The method is applicable to solutions which are free from appreciable quantities of iron Or other interfering ions.
Literature Cited (1) Congdon, L. (2)
*.,and
J., Chem.
128, 98 (1924).
craig, T.J. I,,J ,sot. Chern,In&30, 184 (1911).
(3) Pavlinova, A.
V., J. Applied Chern. (U. S. S.E.), 9, NO. 9,
1682-
9 (1936).
(4)
**
H*i
J*
Chern*
249
457 (lgo2).
R~~~~~~~September 10, 1938. Contribution NO. 52 from the Chemioal Laboratories of the University of LTtah.
Pycnornetric Determination of Lead as Sulfate W. WALKER RUSSELL AND J. H. A. HARLEY, JR. Brown University, Providence, R. I.
The new method in pycnometric analysis has been adapted to the determination of lead as sulfate in two nonferrous alloys with satisfactory results. The behavior of the lead sulfate precipitate is such that centrifugalization can be eliminated, thereby considerably simplifying the method.
I
N A RECENT preliminary report (3)upon a new method in pycnometric analysis, its underlying theory and probable accuracy were discussed, and pycnometric analyses were presented for the barium, iron, or silver in several simple substances, primarily salts. The present work deals with practical analyses of two nonferrous alloys of widely different lead contents. The lead was determined as sulfate because this is the compound most frequently used for its separation and gravimetric determination. The relatively high density of lead sulfate is an asset in the pycnometric method, while its solubility offered an opportunity to apply the pycnometric method to a precipitate of considerably greater solubility than any studied previously in this laboratory.
Density of Lead Sulfate Inasmuch as a pycnometric analysis actually determines the volume of a precipitate, the weight can be ascertained only if the density of the precipitate, or strictly speaking the apparent density characteristic of the experimental conditions existing during the analysis, is accurately known. Other things being equal, the more soluble the precipitate, the greater will be the difference between its absolute and apparent density. I n order to determine the apparent density of lead sulfate, a primary standard of accurately known lead content must be subjected to the pycnometric procedure which is to be used subsequently in the analysis. The principal primary standard used in the present work was the Bureau of Standards lead block, No. 49, melting point 327.3' C. The results of nine determinations, leading to an average value of 5.791, are given in Table I. This value is to be used in computations of analyses in which 3 per cent sulfuric acid is used as the standard washing medium in approximately the amount and manner specified below. The values for the density of lead sulfate given in the literature show a considerable variation. The numerous values listed in Mellor (2) vary from 5.97 to 6.393, while the value given in the International Critical Tables is 6.2. Furthermore, Krings (I) found that lead sulfate formed in
aqueous solution has a lower density (6.03), than that formed by fuming with sulfuric acid (6.272), and that ignited lead sulfate has a lower density than the freshly prepared precipitate. A comparison of the apparent with the real density of lead sulfate therefore required a determination of the latter value for the lead sulfate precipitate formed under the conditions of the analyses. The lead sulfate precipitates in the present work redissolved in the strong sulfuric acid during the fuming and were formed again upon dilution with water. I n six density determinations values of 6.309, 6.277, 6.290, 6.298, 6.338, and 6.292 were obtained, leading to a mean value of 6.301 for the real density of the lead sulfate precipitate. Thus, the real density of the lead sulfate precipitates in the present work is 6.301 while the apparent density when using 3 per cent sulfuric acid is 5.791. The densities of the sulfuric acid used in the present work varied between 1.02475 and 1.0351 a t 30' C.
Apparatus The apparatus employed was essentially the same as that already described (3). Most of the work has been carried out with 1-cc. pycnometers, although 3-co. pycnometers were occasionally used. In the latter part of the work the hand-made pycnometers were replaced by vessels of similar design provided with precision ground-glass joints of standard taper (made by the Scientific Glass Apparatus Company, - - . Bloomfield, N. J.) which have proved very saiiifactory. The evaooration of liauid from the orecioitation flasks can be efficienily conducted- by surrounding ihe vessel with a close-fitting conical asbestos pa er hood, which is readily molded from wet asbestos paper and folds its shape after baking. It is held in position around the flask by a separate, split collar of asbestos paper which is kept firmly in place around the narrow flask neck by a twist of wire. These hoods are readily slipped on and off the precipitation flasks as required.
Method The analytical procedure is in general similar to that already described (3). Since the alloys analyzed contain tin
TABLEI. APPARENT DENSITY OF LEADSULFATB Weight of Lead Block 49 GTam
0.20735 0.2079 0.20835 0.20905 0.2002 0.499s 0,50045 0.19945 0.19985
Apparent Density at 30' Grame/oc. 5.746 5.827 5.771 5.832 5.807 5.77s
5.785 5.778 5.791 Av. 5.791
