dry ice-acetone bath and filtering. The ampoules were m r m e d to room temperature, mixed thoroughly, then cooled s l o ~ l y and , the temperature a t which a turbidity was first observed was recorded. A Dewar flask, a thermometer, some acetone, and dry ice were used t o cool the sample and observe the turbidity temperature. The observations are shown in Figures 1 to 4. As anticipated from Equation 4, straight-line plots are obtained in Figures 1 and 3. Of interest, however, is the indication of two straight lines in each plot. This probably arises from a change in the turbid phase from ice to water. The heats of solution for these two phases would be different, causing a change in s!ope of the straight line. The temperature a t which the change occurs indicates that
the turbid phase is not pure water but mercaptan in water. To make the curves more readily usable for determining tvater in the mercaptans, they are plotted in more convenient units in Figures 2 and 4. The turbidity method for determining ivater in methanethiol and ethanethiol is not intended to be extremely accurate. The turbidity temperature was determined with an accuracy within about 3’ C., which is equivalent to 0.027, water a t a concentration of 0.2 weight 70. An ever-present source of error is the presence of foreign substances in the mercaptan sample, which can give rise to turbidity. On the other hand, the method is very rapid and simple and has been used in these laboratories for several years.
ACKNOWLEDGMENT
The author wishes to express his appreciation to D. 0. Alford and A. B. Menefee for assistance in obtaining infrared spectrophotometric data, and to H. E. Howard for mass spectrometric data. Permission of the Union Oil Co. of California to present and publish this paper is gratefully acknowledged. LITERATURE CITED
(1) Mitchell, J., Jr., and Smith, D. bI.,
“Aquametry. Application of the Karl Fischer Reagent to Quantita-
tive Analyses Involving Water,”
pp. 134 ff, Interecience, New York, 1948.
RECEIVED for review June 11, 1956. Accepted December 4, 1956.
Other papers presented in the group session on analytical research will be published in the March issue of ANALYTICAL CHEMISTRY
Thermogravimetric Analysis of Complex Mixtures of Hydrates EDWARD J. GRIFFITH Research Department, Inorganic Chemicals Division, Monsanfo Chemical
b
The Chevenard thermobalance has been used to determine the quantity of water present in each phase of complex mixtures of anhydrous salts and hydrates, containing as many as six phases. The method is based on the fact that when the proper rate of heating is employed, a selective decomposition of the phase with the highest dissociation pressure occurs. When the phase of the highest dissociation pressure is completely decomposed, the substance with the second highest dissociation pressure begins to decompose, etc. The water present in each phase may b e determined with an average error of about 1 %.
M
ANY OF THE SOLID CHEMICALS
used in household and industrial applications are complex mixtures of hydrates. Common examples are fertilizers, detergents, pharmaceuticals, and industrial poisons. The physical and chemical properties of these solids often depend upon the quantity and location of water of hydration in the products. 198
ANALYTICAL CHEMISTRY
Co., Dayton,
S o satisfactory method has been available for quantitatively determining the water in the various phases of partially hydrated substances, as the water moves toward an equilibrium condition. Simple crystalline mixtures of anhydrous salts and their hydrates may be quantitatively analyzed by x-ray analysis. K h e n the mixture becomes complex, the x-ray method is inadequate and phases are often missed. It is sometimes desirable to know not only how much water is present in a mixture of hydrates and in which phase the water resides, but, even more important, how much water is present in each phase. At any fixed temperature, a system of anhydrous salts and their hydrates exhibits a definite pressure of water vapor, sometimes referred to as the dissociation pressure. If the vapor is slowly removed, the hydrate maintains the same dissociation pressure until that phase is completely dehydrated. If a lower hydrate exists, the pressure drops to the equilibrium pressure of this phase when the higher hydrate disappears. The equilibrium between hy-
Ohio
drates and the rapor phase a t constant temperatures is well known in phase studies. K i t h the development of the modern thermobalance ( I ) , a tool i s available in which temperature may be varied while the atmospheric partial pressure of water vapor over the sample is nearly constant. The dehydration is then followed by measuring the loss of weight as a function of time. The rate a t which a hydrate loses water in an open system at a fixed temperature is dependent upon its dissociation pressure. Thus, the water in two different phases of a mixture of hydrates may be observed by the rate of loss of water from the mixture ( 2 ) . EXPERIMENTAL
The thermobalance employed in this work was manufactured by the A.D.A.M.E.L. Co., Paris, France, and was equipped with a strip recorder obtained from the same company. The temperature of the furnace could be automatically marked upon the pyrolysis curves by an especially designed shorting circuit across two terminals of the
teen 1-ohm standard resistors; and, when activated, one resistor was removed from the circuit of the bucking potential, thus causing the meter relay to return t o zero. The bucking potential was maintained with a carbon cell and could be adjusted to any desired value by adjusting the potentiometer across the leads of the cell. The potential drop across the resistors was calibrated with a hand potentiometer before each determination.
recorder body. The circuit used to perform this function is found in Figure 1.
The furnace temperature was measured with a platinum-platinum-rhodium thermocouple. The furnace therniocouple was bucked with a second thermocouple, maintained a t a measured temperature near room temperature and insulated to prevent sudden change in its temperature. The difference in potential of the therniocouples was fed into a millivoltmeter relay. K h e n the points of the meter relay made contact, an electronic relay was activated which simultaneously activated a time-delay relay and rotary solenoid. The timedelay relay was normally open and shorted the recorder for 30 seconds when activated, causing the pen of the recorder to move down scale. The rotary solenoid was fitted with seven-
Thus, with a knowledge of the room temperature and the potentials a t which the meter relay would make contact, 16 temperatures could be marked on each pyrolysis curve and the temperature range could be varied as desired. The precision of the temperature measurements was about &3' in most instances, but varied somewhat with the temperature range employed. The balance was adjusted to a sensitivity of 5.2 em. per 100 mg.. and a total load change of approximately 400 mg. was usually observed. The chemicals used in the work nere either analytical reagent grades or technical grades n-hich 11-ere purified by n-ell knon-n, reliable methods. A knowledge of the dissociation pressures of the hydrate systems to be analyzed as a function of temperature is an aid in determining the temperature schedule of an analysis, but is not absolutely necessary The usual procedure employed was to dehydrate each component of a mixture separately a t some given temperature schedule before attempting an analysis of the mixture. With the above information a t hand, mixtures were prepared in such ratios that there could be no mistake in the sequence of dehydration-about 25% of the n-atpr !vas present in one phase while 75% resided in a second phase, etc. The lolvest possible maximum temperature treatments which did not
Figure 1. Block diagram of temperature-marking circuit A.
Recorder Time-delay r e l a y Rotary solenoid Rectifier Electronic relay F. Meter relay G , Room-temperature thermocouple H. Furnace thermocouple 1. Carbon cell 6. C. D. E.
0:
allox excessively slow rates of dehydration were used throughout the work. It was found that the lower temperatures yielded sharper breaks in the pyrolysis curves and also gave more readily observed indications of the steps in the loss of mater from a given phase. The exact temperature in an analysis of hydrate mixtures is not usually important; and differences as great as 30" do not markedly influence the final results in most cases, although in some instances where the dissociation pressures of the salts of the system do not differ appreciably, temperature control is critical. RESULTS
Sodium tripolyphosphate (Na5P,010) forms a hexahydrate when crystallized from aqueous solution by the addition of methanol, or acetone, to the solution. The hydrate may be destroyed by heating it near 100" C.; the final product is not anhydrous tripolyphosphate, but a complex mixture of ortho- and pyrophosphates (3). Only five of the six n-ater molecules associated with the hexahydrate are lost during dehydration in which the maximum temperature does not exceed 165'. The one remaining water molecule has reacted with the tripolyphosphate to degradate it t o ortho- and pyrophosphates. Khen the mixture of ortho- and pyrophosphates is heated for a long period, at temperatures of 165' or abore, the sixth water molecule is lost. Figure 2 is a pyrolysis curt-e of the dehydration of sodium tripolyphosphate hexahydrate. There is always an induction time in the dehydration of sodium tripolyphosphate hexahydrate, but the loss of water during this induction time does not correspond to a simple stoichiometric quantity. There is a change in the slope of the pyrolysis curve after about 0.3
\ \
1
0
I
2
3
4
5 6 Hl"6
7
8
9
10
11
0
I
HOYCS
1 Hours
Pyrolysis curves Figure 2. NafiP:O10.6H20 temperature of 100' C.
at
a
constant
Figure 3. Na4Pz07.10HzO temperature of 100' C.
at
a
constant
Figure 4. Na4P207.10H20 Na~Pa010.6H20
in
presence
VOL. 29, NO. 2, FEBRUARY 1957
OF
199
molecule of water has been evolved per molecule of tripolyphosphate. The common hydrate of tetrasodium pyrophosphate (Na4Pz07)is the decahydrate. If a sample of sodium pyrophosphate is dehydrated at a minimum temperature as low as loo", all ten waters of crystallization are lost. Figure 3 is a pyrolysis curve of the dehydration of pyrophosphate. It may be observed from Figures 2 and 3 that the rate of dehydration of pyrophosphate is much greater than the rate of dehydration of tripolyphosphate a t a constant temperature of 100". This leads to the conclusion that one may differentiate between the water evolved from pyrophosphate and that evolved from triphosphate when both hydrates are contained within a single mixture, provided the two hydrates do not decompose simultaneously. Figure 4 is a pyrolysis curve of sodium pyrophosphate decahydrate contained in sodium tripolyphosphate hexahydrate. The interest in this curve is focused on the pyrophosphate. If it is possible to determine the quantity of water contained in one of the phases, the quantity of water contained in the second phase may be obtained from the total loss of m t e r . I n this instance, 200 mg. of water were added as the pyrophosphate decahydrate and 196 mg. were recovered. The water in the tripolyphosphate was ignored, as it did not interfere with the pyrophosphate analysis. Some experience is required to locate correctly the midpoint of the breaks in the curves. I n general, the loner the temperature a t which a break is approached, the sharper the break becomes, and the first noticeable change of slope corresponds to the stoichiometry.
When faster rates of heating are employed, the stoichiometric point will tend to move upward in the break in the curve. This seems to be caused by a slightly premature loss of water from the phase with the second highest dissociation pressure. The curve shown in Figure 4 is a n extreme example, because the time scale was intentionally expanded to enlarge the area of the break in the curve. The stoichiometric point in this curve was obtained at the point of divergence of the tangent to the straight, horizontal portion of the curve to the curving approach. A comparison of the break on the expanded time scale with the other examples illustrates the added ease in fixing the stoichiometric point of the breaks when a slower chart speed is employed for the time coordinate. A pyrolysis curve of a mixture of sodium carbonate monohydrate and sodium pyrophosphate decahydrate at a constant temperature of 100" is shown in Figure 5 . There is no indication on the pyrolysis curve of where the pyrophosphate decahydrate n as completely dehydrated or where the sodium carbonatemonohydrate began to dehydrate. The dissociation pressures of sodium carbonate monohydrate and tetrasodium pyrophosphate decahydrate are very nearly the same a t higher temperatures, but at lower temperatures (60" and below) the differences in dissociation pressures become appreciable. If a mixture of sodium carbonate monohydrate and pyrophosphate decahydrate is decomposed a t a constant temperature of BO0, there is a distinct change of slope a t the stoichiometric point of the pyrolysis curve. The break in the curve beconies even more
pronounced if the temperature of the furnace is allowed to rise slonly from room temperature to 100". I n Figure 6 the pyrolysis curve of the same mixture is shoivn, in which the temperature of the furnace was allowed to rise from room temperature near 30" to 100" over a n 8-hour period. This demonstrates the need for a proper choice of a temperature schedule. I n the analyses discussed to this point, the only anhydrous phases present were those resulting from the decomposition of a hydrated phase. Therefore, at no time iyas the anhydrous salt of a substance of lower decomposition pressure in contact with the hydrated phase of a second salt of higher decomposition pressure. The question arose as to whether the water of the salt of higher decomposition pressure ivas quantitatively lost to the atmosphere or if part of the Fater moved into the anhydrous phases of the salts with loJver decomposition pressures. Figure 7 s h o w the pyrolysis curve of a mixture of sodium carbonate, sodium carbonate monohydrate, sodium pyrophosphate, sodium pyrophosphate decahydrate, sodium tripolyphosphate, and sodium tripolyphosphate hexahydrate. The presence of anhydrous phases does not interfere nith the analysis, nor does it change the shape of the pyrolysis curve or the rate of dehydration of the hydrated substances to any noticeable extent. The same analytical precision was found in both the absence and presence of anhydrous phases. Figure 8 is a typical dehydration of sodium tetraborate decahydrate with a temperature schedule from room temperature to 500" in 8 hours. The hydrate decomposes in steps. The formation of the pentahydrate is clearly
I 0 -
I
Hours
0
I
2
3
4
5
6
7
8
9
1011
12
0
1
2
3
4 H o5u r = 5
7
8
9
HOW6
Pyrolysis curves Figure 5. Mixture of NarP2O.l. 10Hz0 and NazC0a.H20 a t 100' C.
200
ANALYTICAL CHEMISTRY
Figure 6. Mixture of NarP2O,. 1 OHzO and Na2C03 .Hz0 with slowly rising temperature
Figure 7. Mixture of NasPaOio, NasP~Oio.6Hz0, Na4P207, Na4Pa07.1OHzO, NanCOa. and NotCOa. H20
i 0
1
2
3
5
4
6
7
8
,
c
9 1 0 1 1
, I
, 2
, 3
I 4
H0U.l
5 6 7 Hours
8
9 1 0 1 1
Pyrolysis curves Na2BI07.1 OH20
Figure 8.
Figure 9. Mixture of Na26aO7.1OH80 and NaaPzO7.1OH20
L
I \
\
N%C03.H20
I
I I
1
I c0c12 I
y 6H2@
' 200- I1
cobalt chloride hexahydrate. I n this case also, the magnesium sulfate tends to lose its water in steps. The formation of magnesium sulfate tetrahydrate is evident from the pyrolysis cur\ e, and the change of slope is more pronounced betn een the t n o hydrates of magnesium sulfate than between the magnesium sulfate and cobalt chloride hexahydrate. The cobalt chloiide hexahytiiate loses only fire of its SIX water molecules undei the conditions of this analysis. Figure 11 is the pyrolysis curve for a misture of sodium carbonate monohydrate and potassium carbonate sesquihydrate. The formation of the half hydrate of potassium carbonate is also evident from the curve. Tulile I is a compilation of the analytical results obtained n ith the thermobalance. These values \rere not necessarily taken from the pyrolysis curl es. The pyrolysis curves are typical evamples of the system studied, nlieiens the data of Table I correspond to specific esaniples. The aveiage erroi in the work is + 170, and the number of constituents that may be determined a t one time seems to be limited only by the c a p x i t y of the thermobalance. Mole complex mivtureq than thoie mentioned have been analyzed I! itli the thermobalance, but because they add little to the over-all picture, they are not discussed a t this time. Any constituent in a mixture of hydrates may be determined n i t h a n average error no greater than il%, provided a t least 20% of the total nater present is in this particular phase. Phases that contain less than 20% of the total water may also be analyzed. but the accuracy is not as good as that nientioned allore.
3001
I DISCUSSION 0
2
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5
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7
8
9
IO
II
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0
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Hours
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HOYrl
Pyrolysis curves Figure 10. Mixture of M g S 0 4 . 7 H z 0 and CoClz. 6 H 2 0
Figure 1 1 . Mixture of Na2COa.H20 and K2C03.11/2 H20
evident. K h e n the sodium tetraborate decahydrate is dehydrated alone, all ten water molecules are lost. Figure 9 is a pyrolysis curve of a mixture of sodium tetraborate decahydrate (SazB407.10H20) and tetrasodium pyrophosphate decahydrate (NaP20710H20). This curve is of interest because of the similarity of the two compounds. The break in the curve no longer corresponds t o sodium tetraborate pentahydrate, and the sodium tetraborate no longer dehydrates in steps but loses its water continuously. Furthermore, only nine molecules of water are lost from the sodium tetra-
borate a t a maximum temperature of 165". At temperatures as high as 500", a part of the tenth n-ater molecule is lost in 8 hours, but appears in the pyrophosphate phase. This problem is easily solved in the two-phase mixture by calculating the water evolved to the first break as only nine tenths of that present in the borate phase and using sufficiently high temperatures to dehydrate both phases completely. Some other approach Ivould be necessary for more comples mixtures, however. In Figure 10 a pyrolysis curve is shown for the dehydration of a mixture of magnesium sulfate heptahydrate and
Throughout this ork the primary concern has been to determine the quantity of water contained in the various phases of partially hydrated mixtures. The method is equally well adapted to the analysis of the crystalline substances nhich contain the water in the mistures, if all phases are completely hydrated prior to their decomposition. For example, the phosphate constituents may be determined in a mixture of ortho-, pyro-, and tripolyphosphates by first adding enough n-:tter t o the mivture to hydrate it completely. The sample is next air-dried, and the pyrolysis curve is obtained for the sample. I n this case, the emphasis is placed upon the quantity of each phosphate present rather than the water contained in the phosphate. Probably the greatest single source of error which may be encountered in the thermogravimetric analysis of hydrates results froin adsorbed nater. I n the VOL. 29, NO. 2, FEBRUARY 1957
201
Table 1.
Hydrates Analyzed NarP207.10H20 NaZCOs.HZ0 NasPsOio. 6H20 Na4P207.10H20 NarPnO,. 10H20 Sa2COj.H20 NajOsOto.6Hz0 plus anhydrous phase of each hydrate YazB40.i. 10H20 SazB107.lOHzO NakPz07. lOHzO Na2C08, HzO KzC03.1.5HzO MgSOd. i"z0 CoClz. 6Ht0 0
Analysis of Mixtures of Hydrates
H20 -Added, Mg. 214 74 96 407 189 77 93 378 216 94 73 245 112 192
Hz0 Found, Mg, 212 73 96 403 189,190,188 74,79,77 90,93,93
Error,
377 216" 95
-0.2
9%
-1
-1.4
0 -1
0 4-1
e13 +2.8 247 +0.8 107 -4.5 193 $0.5 Absolute average 1. 2 22 mg. of HzO from -ha2B4O7.10HzO phase found in Sa4Pz07phase.
work reported here, the adsorbed moisture did not influence the results t o a detectable degree, but i t is conceivable that with mixtures in which the various
hIax. Temp after 8 Hours, "C.
165 165 100 165
500 500 165 165
phases were completely hydrated adsorbed n-ater might contribute to the total weight loss of the sample. It is believed that when anhydrous phases
are present, adsorbed water is converted to water of crystallization and does not contribute an error in the analyses. Nevertheless, it is advisable t o be continuously alert to this possible source of error. ACKNOWLEDGMENT
The author wishes to express his gratitude to John Andersen of the Monsanto Chemical Co. for designing and building the temperature-measuring unit used in this work. LITERATURE CITED
(1) Chevenard, P., Woche, X., De la Tullayl, R., Bull. soc. chim. 5 ,
10 (1944). (2) Oroeco, E., Ministerio trabalho ind. e corn. Inst. nacl. technol. (Rio de Janeiro) 1940,33. (3) Quimby, 0. T., J. Phys. Chem. 58,603 (1954). RECEIVEDfor review July 19, 1956 Accepted October 5, 1956. Divisions of .Analytical and Physical and Inorganic
~~~~~~v~~~~~~~~~~~~~~~~~~~~~
129th hleeting, ACS, ~ ~ lT ~l ~~ , ,~~ , ~ 1956.
Phosphori metry A New Method of Analysis R. J. KEIRS, R. D. BRITT, Jr.,l and W. E. WENTWORTH Deparfmenf of Chemisfry, Florida Sfafe Universify, Tallahassee, Flu.
bPhosphorimetry is a means of chemical analysis based upon the nature and intensity of the phosphorescent light emitted by an appropriately excited molecule. Many organic mol: ecules containing multiple bonds, when in a rigid glass formed a t low temperature by solutions of the material in suitable solvents, phosphoresce if excited by radiant energy of suitable frequency. Each phosphorescence is unique in regard to its frequency, lifetime, quantum yield, and vibrational pattern and such properties are used for qualitative identification. The correlation of intensity with concentration can serve as a basis for quantitative measurement. Mixtures are analyzed by the use of a resolution phosphoroscope. The method has been applied to the determination of several compounds and their mixtures, three of which are described.
Present address, Savannah River Plant, E. I. du Pont de Xemours & Co., Inc., Augusta, Ga.
202
*
ANALYTICAL CHEMISTRY
T
HE
PHOSPHORESCESCE
EMISSIONS
from over 200 compounds have been reported by molecular spectroscopists interested in the elucidation of the energy schemes of molecules. About 90 such emissions have been tabulated by Lewis and Kasha (11). The use of phosphorescence emission spectra for identification of substances was first suggested by Lewis and Kasha; however, this subject has been developed very little in the past decade. More recently the idea of using a resolution phosphoroscope (proposed by M. Ilasha of this laboratory) greatly extended the annlytical potentialities of phosphorescence. B theoretical analysis of the method is to be published. Phosphorescence is not characteristic of a specific class of compounds, but a prime condition for its observation is high viscosity. Substances that phosphoresce may be divided into two classes, based upon the mechanism by which their phosphorescences are produced (11). I n the first group are
mineral, or crystal, phosphors (15). I n this case the individual molecule is not phosphorescent, but the ability to emit a n afterglow is associated with the return of an electron to a n impurity site in the crystal, following ionization through the process of photon absorption. -4s this type of phosphorescence cannot be ascribed to a definite substance, phosphorescences of this clash were not considered in this investigation. I n the second class the emission is attributed to a definite molecular species. whether the substance is in a pure crystalline state, absorbed on a suitable surface, or dissolved in a suitable rigid solution. It is with such rigid transparent solutions that this investigation was concerned. Phosphorescence is produced in molecules by the absorption of radiant energy of a frequency within the normal absorption band of the molecule-usually in the ultraviolet region of the spectrum. Two of the many possible end results are fluorescence and phosphor-
i
