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INDUSTRIAL AND ENGINEERING CHEMISTRY
Table I11 gives the average results obtained when various quantities of lime are added up to the point where the precipitation of sugar begins. In making the calculation a correction was made, as usual, for heat of hydration of the excess calcium oxide.
VOL. 30, NO. 11 b . .
This discussion has touched on a number of the problems involved in the industrial application of the reactions between calcium oxide and sucrose. Many points of theoretical interest remain to be explored-for example, the actual chemical composition of the soluble phase, which varies with initial TABLE111. HEATOF REACTION AT VARIOUS STAGES OF FORMA- sugar concentration; whether the precipitate actually is TION OF SOLUBLE PRODUCTS tricalcium saccharate or a mixture of compounds; what sort CaO per Mole Heat per Mole CaO per Mole of energy relation exists which favors the occurrence of a Sugar Added Sugar in S o h . CaO reaction that gives 12.5 Calories (the precipitation of triMoles Moles Cal. calcium saccharate) in preference to one that gives 16 Calories 22-25 21.1 1.09 1.00 (the hydration of lime). 20.0 1.47 1.38
19.6 2.01 1.63 19.0 2.55 1.90 2.60 * Precipitation begins 0 Results with this and smaller quantities of lime (less than 10 grams er cooler) were erratic and uncertain. Larger amounts gave reproducigle results.
The heat of reaction during the period of precipitation still remains to be determined. Direct determination of the total heat generated during this phase of the reaction was made, and the heat due to the reaction itself was calculated by correcting for the heat of hydration of the excess lime. This method gave an average value of 12.7 Calories per mole of calcium oxide. An indirect calculation based on 16 Calories per mole of calcium oxide for the complete reaction and 19 Calories for the solution-phase reaction, and using actual experimental values for sugar precipitated, sugar and calcium oxide in solution, etc., gave a value of 12.2 Calories. Calculations from the experimental data indicate that the most probable value for heat of formation of tricalcium saccharate from solid calcium oxide in 5.0 to 5.5 per cent sugar solutions a t 8” to 20” C. is 46.5 Calories per mole of saccharate, which is equivalent to 15.5 Calories per mole of calcium oxide. T o summarize, the heat of reaction during the formation of the soluble phase is 19.0 Calories per mole of calcium oxide, during the precipitation phase 12.5 Calories per mole, and for the completed reaction (including the residual alkalinity) 16.0 Calories per mole. Additional heat is evolved as a result of the partial hydration of the excess lime added. This amounts to about 6.0 Calories per mole of excess calcium oxide. Variations in the quantity of total heat are due to this latter source which is dependent on the amount of excess lime, the temperature, and the time of lime addition. This accounts for the variations in total heat shown in Tables I and 11. Von Lippmann,’ reporting the work of Pettit, gives the heat of formation of the monosaccharate from solid calcium hydroxide in sugar solution a t 70” C. as +7.2 Calories per mole, and of the disaccharate under similar conditions as +11.7 Calories per mole of saccharate. One mole of monosaccharate and one mole of solid calcium hydroxide gives +4.5 Calories. Pettit also says that the further addition of sugar solution to a solution of calcium monosaccharate gives an evolution of heat amounting to a maximum of 3.1 Calories. Concentration, temperature, and other conditions are not given. This statement is in agreement with the writer’s observation that the heat of reaction is greater at the beginning of the period of lime addition-that is, when the ratio of sugar to calcium oxide is greatest (Table 111). Herdeldl gives the heat of reaction between 1.23 grams of caustic lime (calcium oxide) and 75 grams of 10 per cent sugar solution a t 18” C. as 8.63 Calories (presumably Calories per mole). This figure is not in agreement with results reported here. Under similar conditions 21.1 Calories per mole were liberated (Table 111). “Chemie der Zuckerarten,” 3rd ed., Part 11, p. 1332, Friedrick Vieweg & S o h , 1904. 1
Acknowledgment The writer wishes to acknowledge his indebtedness to Robert J. Brown, A. N. Bennett, and others for much of the experimental data which served as a basis for this paper. RECEIVED April 25, 1938. Presented before the Division of Sugar Chemistry at the 95th Meeting of the American Chemical Society, Dallaa, Texas. April 18 to 22, 1938.
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CORRESPONDENCE
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Chemical Industry and Economic Progress SIR: In Table X, page 941, of the article entitled “Chemical Industry and Economic Progress” [IND.ENQ.CHEM.,30, 93442 (1938)], the price change for anhydrous ammonia was given as an increase of 15.6 per cent. This change applies to cylinder ammonia which was the article used in the Chemical and Metallurgical index on which the study was based. The following series of tank car ammonia prices was obtained from Chaplin Tyler of the du Pont Company since prices for the years 1926 to 1929, inclusive, were not published in the leading chemical journals. The price is not available for 1925 because anhydrous ammonia was not then available in tank cars: Year 1926 1927 1928 1929
Cents/lb. 7.5 7.0 6.25 5.50
Year 1930 1931 1932 1933
Cents/lb. 5.50 5.50 5.50 4.50
Year 1934 1935 1936 1937
Centdlb. 4.50 4.50 4.50 4.60
From these figures it may be determined that the average price for the four years 1926 to 1929, inclusive, was 6.56 cents per pound, whereas it was 4.50 cents per pound for the four years 1934 to 1937, inclusive. The progress ratio computed as the ratio of the more recent four-year average t o the earlier average is therefore 68.6 per cent, corresponding t o a price reduction of 31.4 per cent. Accordingly, if tank car ammonia prices had been used in place of those for cylinder ammonia, a price reduction of 31.4 per cent would have been obtained instead of an increase of 15.6 per cent. This would have placed anhydrous ammonia well up in the group of so-called developmental chemicals instead of placing it at the bottom of the group of so-called stabilized chemicals. D. P. MORGAN SCUDDER. STEVENS & CLARK Nmw YORK.N. Y. September 8, 1938
