August, 1928
INDUSTRIAL A N D ENGINEERING CHEMISTRY
Effect of Storage Period on the Jelly Strength
Jellies increased in strength with length of storage, as shown by Figure 3. A storage temperature of 23" to 25" C. was used and all jellies were paraffined within a few minutes after filling the glasses. The maximum change occurred very soon after making, but a gradual increase took place over a period of several days. This confirms Tam's6 work. It is recommended that jellies be allowed to stand a t least 24 hours before testing. Weak jellies were found to set more slowly than firmer ones, but the changes taking place after 24 hours were usually slight. Contrary to Tarr,G it was found that the surface of uncovered jelly toughened very noticeably on standing. This skin effect or surface hardening became evident as soon as the jelly cooled and increased as long as the jellies remained uncovered. I n order that an accurate determination could be made it was found necessary to cover the surface of the jellies with melted paraffin or a hermetic seal soon after making. To illustrate, a current jelly was prepared in the regular 2-ounce (57-gram) jelly glasses. One-half of these were paraffined at once while the rest remained uncovered. After 4l/2 hours the paraffined jelly tested 93 grams while the uncovered tested 147 grams. After 52 hours' storage at 23" C.
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the readings were, respectively, 131 and 174. In every case uncovered jellies showed this surface hardening effect. Limits of Accuracy of Instrument
Table I shows the results obtained from a large number of determinations. The probable error or the mean of any series of determinations was found in all cases to be within - 3 grams. This was calculated according to the formula:
4%
-
P.
E.
= 0.6745
where P. E. is probable error of the mean, Zd2 is equal to the summation of the squares of the individual deviations from the arithmetical mean, and n equals the number of determinations. Considering both top and bottom readings in Table I, the average per cent deviation of tn.enty-two series of samples was approximately 5 per cent. There was little difference in accuracy between series run on the bottom or on the top of jellies. The jelly strength of the bottoms of the samples of apple jellies and cranberry sauces were usually less than those of the tops. The reverse was found to be the raw with artificial pectin jellies. For practical purposes it is not considered necessary to test the bottom of jelly samples.
Note Concerning Thermodynamic Calculations' David F. Smith P I T T S B U R G H EXPERIMEICT S T A T I O N ,
u. S.B C R E A V
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STEREST reeelilly attaching to the therniodynamic properties of hydrocarbons and alcohols makes it desirable to attempt to clear up some apparent misunderstandings which the writer's paper2 on the subject seems to have caused. Two interesting art'icles by Francis3 have since appeared in which the free energies of various hydrocarbons and alcohols have been estimated. The present writer wishes to commend these efforts to present the best available thermodynamic data, uncertain though they may be, in a form that makes them usable for those who would not take the .trouble to search the literature, make the calculations, and correlate the data for themselves. Although many of the data that Francis uses are admittedly but rough approximations, his free-energy equations, especially in cases like CH4, C2H2,and C2H4,in which some equilibrium data are available, do represent in condensed and usable form such probable thermodynamic properties as are knonm. If a free-energy equation is consistent with such experiment'al work as has been done, it is a very concise and convenient representat'ion of the known facts. However, even when one has reliable thermodynamic data, the story is but inconipletely told. Although the thermodynamic facts cannot be violated, after all, the rate of reaction is ultimately the determining factor for practical purposes. For example, Francis3 thinks that the chance of forming benzene f:om water gas, although thermodynamically quite possible, is small in comparison with that of forming the more stable hydrocarbons and alcohols. However, in spite of the fact that benzene is highly unstable thermodynamically with respect to decomposition into its elements, it is nevertheless a major product from high-temperature distillation processes. Again, alt'hough thermodynamica!ly Received April 30, 1928. Published by permission of t h e Director, U. S.Bureau of Mines. ( N o t subject t o copyright,) I N D . EKG. C H B M . , 19, 801 (1927). 8 I b t d . . 2 0 , 2 7 7 , 283 (1928).
OF hfINES, PITTSBURGH,
P.4.
more stable at lower temperature nith respect to such decomposition, benzene does not appear in lowtemperature coal tar, In the Fischer process for synthesis of hydrocarbons from water gas, it has been shown4that ethylene is produced in quantities quite comparable to those of ethane, which is far more stable with respect to decomposition into carbon and hydrogen at the temperatures at which this process is carried out. In these cases we see that the rate of reaction is an important factor. Thus the fact that a certain hydrocarbon is far less stable with respect to decomposition into its elements than is another does not necessarily prevent its predominance when formed from other than its elements, for example, from carbon monoxide and hydrogen, as in this case. Although the writer is not recommending that anyone attempt to produce benzene from uater gas, such a process does not seem impossible. As far as requiring six carbon atoms (at least 6CO) for its synthesis from water gas, it is known that higher hydrocarbons (containing twenty or more carbon atoms per molecule), which have a great thermodynamic tendency to decompose into their elements or into lower hydrocarbons, are readily formed from water gas. The thermodynamic tendency to form a given hydrocarbon from carbon monoxide and hydrogen is in general quite different from that to form the hydrocarbon from its elements. If the rate of forming the hydrocarbon from water gas is rapid compared to the rate at which it decomposes into its elements under the same conditions, its formation in this way is quite feasible, irrespective of any considerations as to the inherent instability of the product. The thermodynamic tendency for a reaction to take place depends in part upon the prevailing conditions of concentration. There is perhaps room for a choice of the conditions to be taken as a standard for comparison. The writer prefers, honever, to refer to the conditions that have Smith Davis, and Reynolds, IRD ENG.CHEM , 20,462 (1928).
INDUSTRIAL AND ENGINEERING CHEMIXTRY
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Vol. 20, No. 8
been fairly generally accepted as standard-that is, where all substances involved are a t the standard pressure (more strictly, fugacity) of 1 atmosphere. To take a particular case, consider the reaction
range considered (0” to 400” (3.). Table I in that paper clearly shows that the temperature coefficients of the different reactions are different. A relation between the equilibrium constants of the different reactions, which holds for one temperature range, may not hold for another. 2CO 5Hz = CiHe 2H20. I n conclusion, a few minor errors in Francis’ paper3 should be pointed out. It is inconsistent to use the present writer’s value for the free energy of ethane, which was derived from represents the maximum work obtainable (free-energy in- his value for methane, unless the tatter value is accepted. crease) in the reversible formation of 1 mol of CzHG at a pres- This, however, makes an error of but about 500 calories, sure of 1 atmosphere when each of the substances involved which, in comparison with the reliability of some of the other is maintained at the beginning and throughout the process data used, is perhaps not important. a t a constant pressure of 1 atmosphere. This free-energy Francis3 erroneously compares his value for methane increase is then a measure of the tendency for this reaction to (gas, 1 atmosphere) with Parks and Kelley’s value for take place. It is true that in practice a process such as this methane (liquid, 1 atmosphere) which was derived from in which the pressures of all the substances involved are Lewis and Randall’s6 value for methane (gas, 1 atmosphere). 1 atmosphere a t the beginning and throughout the process The vapor pressure of (hypothetical) liquid methane a t would never be encountered. But this, nevertheless, is the 25” C. and 1 atmosphere would, of course, not be 1 atprocess the writer has chosen to consider as a standard for mosphere. Francis’ statement that “practical pressures will comparison, and this is the general practice in thermo- not usually overcome an unfavorable free-energy change of dynamic considerations. Thus, when the value of the con- more than 5000 calories” is rather loose. The effect of presstant of the reaction, K , is larger, the thermodynamic teiid- sure depends upon the magnitude of the decrease in volume ency for the reaction to take place is greater. And, although accompanying the reaction. As a matter of fact, the synin forming a mol of one hydrocarbon the molal proportions thesis of methanol from carbon monoxide and hydrogen is of reactants and resultants required are different from those carried on a t temperatures where the free energy increase at required for another hydrocarbon, all the free-energy quan- standard pressure is over 9000 calorie^.^ It must be admitted tities that have been compared refer to the formation of 1 -and this is a criticism that applies to certain of the writer’s mol of hydrocarbon from the same reactants and resultants own calculations as well 8s to those of Francis-that, as has a t the same partial pressure. The actual processes corre- been shown frequently in the past and as will be shown sponding to given free-energy changes are set forth especially again in a paper soon to appear from this laboratory, freeclearly by Noyes and SherrX5 energy values calculated on the basis of rough approximaThe statement in the first conclusion of the writer’s paper2 tions to and extrapolations of thermal data are often quite concerning the relative tendencies for different hydrocarbons erroneous. to form from water gas, of course, refers to the temperature 6 “Thermodynamics,” McGraw-Hill Book Co., New York, 1923.
+
6
+
“Chemical Principles,” The Macmillan Co , New York, 1922.
7
Lewis and Frolich, IND.END.CHEX.,20, 285 (1928).
Use of Buffers in the Determination of Color b y Means of Titanium Trichloride’ I-Amaranth, Ponceau 3R, and Orange I 0. L. Evenson and D. T, McCutchen COLOR CERTIFICATION LABORATORY, FOOD,DRUG,AND INSECTICIDE ADMINISTRATION, DEPARTMENT OF AGRICULTURE, WASHINGTON. D
HE original method for the evaluation of color in azo dyes by means of titanium trichloride was formulated by Knecht and HibberL2 The hot-water solution of the dye was reduced in the presence of hydrochloric acid, Rochelle salts, or sodium bitartrate, in an atmosphere of carbon dioxide. For some time this laboratory has carried out the titration of Amaranth (Colour Index 184) in the presence of hydrochloric acid. This method was found objectionable, however, for three reasons: (1) The end point is not sharp; (2) results tend to be slightly low, owing, in small part rtt least, to fading of the dye in acid solution; (3) it does not permit accurate or convenient evaluation of total color in mixtures of Amaranth with other dyes which cannot, or can only with difliculty, be titrated in acid solution. In an attempt to obviate the necessity of using acid, the titration of Amaranth in the presence of various buffers (or catalyst-buffers) was
T
Presented before the Division of Dye Chemistry at the 75th Meeting of the American Chemical Society, St. Louis, Mo., April 16 to 19, 1928. 4 “New Reduction Methods in Volumetric Analysis,” Longmans, Green and Co., New York, 1918. 1
c.
investigated. On obtaining good results in these determinations, the experiments were extended to other dyes. Materials
Commercial samples of dye of certified grade were used, the analysis being made according to the methods prescribed for color ~ertification.~ The following salts were tried as buffers: Rochelle salt, sodium bitartrate, calcium carbonate, sodium acetate, sodium tartrate, sodium citrate, potassium antimony tartrate, and sodium bicarbonate. These were of commercial c . P. grade as obtained from several sources. The 0.1 N titanium trichloride used was standardized with a ferric sulfate solution made from pure ingot iron.a This result was checked by standardization against sodium ~ x a l a t e . ~The titanium solution was 1.4 N with respect to the hydrochloric acid present. Procedure
The buffer salt, in the quantity to be used, was dissolved in water by heating to boiling in a wide-mouth Erlenmeyer 8
U. S. Dept. Agr., Bull. 1390, revised (1928).
