I Figure 4.
Electrolyte frame
I 39'
NGTE: ALL
RIBEIIN~ IS I * T H ! C & M ~ T E R II S~ L C a s r IRON
Acknowledgment
Figure 5.
The development group consisted of Robert Halfon, chemical engineer; Allyn Heit, chemist; Daniel Bauer, junior mechanical engineer; and Walter Dzingala, chemical engineer, with assistance on test runs by Robert Casciano, chemical engineer. literature Cited
(1) Conway, H. F., Sohns, V. E., IND.ENG.CHEM.51, 637 (1959). (2) Dvorch, W., Mehltretter, C. L. (to L.S. A., Secretary of Agriculture), U. S. Patent 2,648,628 (Aug. 11, 1933). (3) Mehltretter, C . L. (to U. S. A., Secretary of Agriculture), Zbid., 2,713,553 (July 19, 1955).
Cell support plate
(4) Zbid., 2,770,589 (Nov. 13, 1956). (5) Zbid., 2,830,941 (April 15, 1958). (6) Mehltretter, C. L., Rankin, J. C., Watson, P. R., IND.ENG. CHEM.49, 350 (1957). (7) Pfeifer, V. F., Sohns, V. E., Conway, H. F., Lancaster, E. B.,Dabic, S., Griffin, E. L., Jr., Zbid., 52, 201 (1960). RECEIVED for review February 17, 1961 ACCEPTEDOctober 18, 1961 Work done under contract with the U. S. Department of .4griculture and authorized by the Research and Marketing Act. Contract supervised by the Northern Utilization Research and Development Division, Agricultural Research Service.
CONTINUOUS AUTOMATIC MANUFACTURE OF HYPOCHLORITE SOLUTIONS FROM SODA ASH ROBERT L. M C B R A Y E R AND N E W L I N S. N I C H O L S Chemical Engineering Research, Wyandotte Chemicals Cor)., Wyandotte, Mich.
The success of continuous automatic systems for the manufacture of hypochlorite solutions from caustic soda and milk of lime led to experimental work on the manufacture of hypochlorite solutions from soda ash. Potential applications for soda ash-based solutions exist in the pulp, textile, and metallurgical industries. Experimental data demonstrated the applicability of oxidation potential as a means of control in a continuous system Reproducible data gave a correlation between oxidation potential and available chlorine concentration. Stability of the solutions as a function of temperature and concentration was studied.
HE USE OF OXIDATIOS POTESTIAL to control the continuous Tautomatic manufacture of sodium hypochlorite solutions was proposed by Pye in 1950 (70). Since that time, information has appeared on experimental work and commercial installations of such systems to produce sodium and calcium hypochlorite (7-3, 8, 72). The sodium hypochlorite solutions have been made from caustic soda. Continuous automatic operation offers two primary advantages over batch operation-reduction of manpower and reduction of required production space. Less evident, but economically important advantages are-a more stable and uniform product and the reduction of the danger of overchlorination.
148
l & E C PROCESS D E S I G N AND DEVELOPMENT
The success of continuous automatic units producing hypochlorites from caustic soda and milk of lime led to experimental work using soda ash. Potential applications for soda ash-based hypochlorite exist in the pulp, textile, and metallurgical industries. In the pulp and textile industries, it may be desirable to use a bleach with a comparatively low p H (between 8 and 9) (73). Soda ash-bleach solutions are buffered to a p H of' 8.5 to 9. In addition, the caustic removal problem is eliminated. For certain metallurgical applications, such as the separation of cobalt and nickel ( 7 7 ) , soda ash-based hypochlorite could supply the chlorination agent and would neutralize the liberated acid.
Soda ash solutions chlorinate faster, more completely, and with less heat generation than do caustic soda solutions. Cooling may generally be eliminated. Economics may not be favorable unless sodium bicarbonate produced during chlorination (Equation 3) can be utilized. Although raw material costs are greater for soda ash systems than for caustic soda systems, the factors given above may shift the economic balance in specific applications.
Control Aspects
I n a n oxidation-reduction (redox) system, a measurable potential difference can exist between two suitable electrodes. For the general case of a redox system given by a-4
+ bB + . . . .
~2xX
+ yY + . . . . + ne
the potential is given ( 4 ) by
Chemical Aspects
Chlorination of soda ash solutions generally yields a solution containing sodium hypochlorite, sodium bicarbonate, excess sodium carbonate, and sodium chloride. The chlorination can be carried to the point where all the carbonate is converted. The reaction proceeds in two steps:
+ C12 NaOCl + NaCl + COZ,AH + 2000 cal. 2NaHC03, AH - 6000 cal. NaSCO, + COS + H d 0
NaSCOj
-f
+
(1) (2)
T h e over-all theoretical reaction, which does not provide for the desired excess sodium carbonate, is C12
+ 2Na2C03+ HSO
+
NaOCl
+ NaCl + 2NaHC03 AH
- 4800 cal.
Conditions can be adjusted so that the potential depends on the concentration of the reacting species. Redox potential changes can then indicate degree of reaction and serve as control references. Redox operating curves are normally developed from experimental data by plotting potential as a function of the concentration of a reactant. If reproducible curves exist, redox can be used to control at desired conditions. Ease of control is indicated by the slope of the curve. Accurate control is difficult where the slope is nearly zero or infinite. I n alkaline hypochlorite solutions, the redox potential is governed ( 9 ) by
(3)
Fisher and Carlson report ( 3 ) that the reaction of chlorine with a n alkali solution requires an induction period, followed by a period of rapid rvaction. This indicates an autocatalytic reaction. Theoretical quantities of pure chemicals required according to Equation 3 are 3.02 pounds of sodium carbonate per pound of chlorine. To reduce the danger of overchlorination (Equation 4), normal practice is to use 3.25 to 3.5 pounds of commercial soda ash per pound of chlorine. Experimental evidence shows that Equation 2 does not always go to completion. Once steady-state conditions for a system were established, the soda ash-chlorine ratio may be reduced to obtain a more active, but less stable, hypochlorite. Soda ash-based hypochlorites are not generally made in excess of 40 gram-per-liter available chlorine because of temperature limitations. For higher available chlorine concentrations, the solubility of the sodium carbonate and bicarbonate would be exceeded.
H20
+ C1*+
NaCl
+ C 0 2 + HOC1
(4)
Hypochlorous acid, under the influence of carbonic acid, decomposes to oxygen and hydrochloric acid (7). The hydrochloric acid reacts with hypochlorite and bicarbonate, and the decomposition continues with increasing violence. T h e temperature of the reacting solution should be kept below 27' C. to prevent formation of sodium chlorate (7) by decomposition: 3NaOCI
-+
2NaC1
+ NaC10,
(5)
Stability of hypochlorite solutions in general is dependent on the following (6, 7):
1. Decomposition catalysts present 2. Hypochlorite concentration 3. Alkalinity of the solution 4. Temperature 5. Exposure to light and air
+ 2 H + + 2e
(7)
Since the potential is primarily a function of hydrogen ion concentration in Equation 8, a sharp inflection in the redox operating curve should occur a t the reaction end point. This point of inflection can be used as a control point. The interdependence of the potential with hydrogen ion concentration and hypochlorite concentration is shown in the case of caustic soda systems by the fact that plotting redox curves for various starting concentrations of sodium hydroxide results in a family of curves ( 3 , 70, 72). Production of hypochlorite solutions from soda ash is a redox system as shown by the following equations. Equation 3 written ionically is
+ H 2 0 F! C10- + CI- + 2HC032C0s-2 + H + F! 2HCO3- neutral
2c03-' f Clz
Stability of Solutions
NaHC03
e C10-
and the voltage is given (70) by
or
If chlorination is continued beyond the total conversion of sodium carbonate, the following reaction occurs :
+ C1-
Clz
2CI- - 2e reduction
(9)
(10)
(11)
and Equation 7 for oxidation. I n alkaline solutions, Equation 8 can be rewritten as E
=
E" -
0.591 [OCI-] [HCOQ-I' log 2 [a-I[CO,-*I
(12)
I n this system, the hydrogen ion concentration effect is minimized, since hydrogen ions in Equation 7 are consumed (Equation 10). T h e solution is strongly buffered. A 0.1 M solution of sodium bicarbonate and carbonate has a p H of approximately 10.2 (5). The redox curve shows a n inflection at the reaction end point, however, the inflection is not as sharp as for a caustic soda system. Materials
Starting materials consisted of chlorine, water, and light soda ash having the following typical analysis : 58.1 99.5 0.02
I n addition to these, hypochlorite made from soda ash is subject to decomposition due to the presence of bicarbonate. Sunder found that 5y9sodium bicarbonate makes the solution very unstable ( 7 ) .
0.32
0.20
14 VOL. 1
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APRIL 1 9 6 2
149
Vent
I
Gas Sample Point
+
n CitfyWater
25 lb.Steam
1
Water
Figure 1. 1.
2.
The continuous automatic manufacture of hypochlorite solutions was studied in this system
One-ton chlorine cylinder Vaporizer
3. 4.
Rupture disk Relief valve
5. 6.
Hypochlorite reactor Electrodes
Liquid chlorine, obtained in 1-ton cylinders, was essentially 1007, pure. T a p water was normally used; however, steam condensate and/or filtered river water were tested with no apparent effect. Apparatus and Procedure
The flow sheet for the experimental work is illustrated in Figure 1. Soda Ash Solutions. Solutions were made up in a 500gallon steel tank. Material requirements are given in Table I. It was necessary to heat the solutions to 20' C. to dissolve the soda ash completely. The solutions were pumped to the reactor by a centrifugal pump. Feed rate was measured by a rotameter. Chlorine Supply. Liquid chlorine was vaporized in a hot water vaporizer maintained a t 50' to 60' C. Gaseous chlorine is preferred to liquid chlorine owing to difficulty in accurately controlling liquid chlorine flow. Chlorine gas passed through a barometric leg made from 1-inch extra heavy black iron pipe to a porous stone filter. The line from the filter to the control valve was 1/4-inch extra-heavy black iron pipe. Chlorine flow to the reactor was controlled by a '/l-inch, G trim, Research Controls valve actuated by a pneumatic signal from the redox recorder-controller. A IO-foot section of line from the valve to the reactor was l/*-inch Uscolite C P (styrene-acrylonitrile-butadiene copolymer) pipe. The chlorine was introduced into the reactor through a n orifice, 3/16 or l/g inch in diameter. Redox Instrument. An Electronik redox recorder-controller (Brown Instrument) was used to measure and record the redox potential of the hypochlorite solution, and to control the process at the desired available chlorine concentration. p H Instrument. A Leeds-Northrupline-operated p H meter was used to measure p H in several runs made with 15gram-per-liter available chlorine hypochlorite solutions. -4 calomel-glass electrode pair was used. Redox Electrodes. A Beckman platinum-silver electrode pair was mounted horizontally through rubber stoppers in the 150
l&EC
PROCESS D E S I G N A N D DEVELOPMENT
7. 8.
Rotameter Soda ash solution tank
9. 10.
Switch to instrument Switch to solution pump
reactor 28 inches above the chlorine inlet. A metal clamp held the electrodes firmly in position. Hypochlorite Reactor. The reactor, shown in Figure 2, consisted of I1/2-inch diameter borosilicate glass pipe supported vertically, with a 3-inch glass tee a t the upper end. The tee acted as a disengaging section for the hypochlorite and any gases present. The lower end of the reactor was a ll/z-inch tee. Soda ash solution entered from the bottom of the tee and the chlorine nozzle entered through the side arm. A I-inch diameter, 2-inch long draft tube, fabricated from Uscolite pipe was centered over the chlorine orifice. This tube set up internal circulation of the hypochlorite, promoting the autocatalytic reaction. The chlorine gas, emerging in small bubbles from the orifice, increases the velocity of the solution through the draft tube, thus setting u p a recirculation current. Valves were installed on the 3-inch tee so that the outflow of gases and hypochlorite could be regulated to provide for pressure operation of the reactor.
Table 1.
Material Requirements for Hypochlorite Solutions Prepared from Soda Ash and Chlorine
Material requirements per 100-gallon hypochlorite solutionG
Auailable Chlorine, G./L. 5 10 15 20 25 30 35 40
Soda Ash Solution, G./L. 'VazCOa 16.5 3311 48.3 66.1 83.3 100.4 117.1 134.4
Chlorine, Lb. 4.2 8.4 12.5 16.7 20.9 25.1 29.2 33.4
Soda Ash (99.5% .VaaZCO3), Lb. 13.7 27.3 40.6 54.3 67.9 81.6 94.9 108.6
Water, Gal. 99.4 9817 98.3 97.6 97.0 96.4 95.9 95.4
Sp. Gr. of Hypochlorite (at 20" C./ 20' C.) 1 015 1.030 1.047 1.061 1.076 1.092 1.108 1.124
Basis: Pounds of chlorine = Gallons hypochlorite solution X g./l. available chlorine 7 19.8 g. gal./lb. 1. Pounds soda ash = pounds chlorine X 3.25. a
Uniformity of Hypochlorite Solutions Made in the Continuous Automatic System Is Excellent Test No. 1 2 Sample period, min. 56 92 No. of samples 11 19 Available chlorine, g./l. 40.6-42.0 15 .O-15.6 Range 41, ? 15.3 Average f0.2 Standard deviation f0.4 Table II.
P
VENT
A S SAMPLE LINE
h A
HYPOCHLORITE
Table 111.
Reproducibility between Runs in the Continuous Automatic Hypochlorite System Initial Redox Available NaZCO3, Potential, Chlorine: Run G /L. MU. G./L. PH 8.7. 1 131 845 40.0 2 131 844 37.3 8.6" 3 131.3 844 42.0 8,6a 4 152 844 40 1 9.lU 6 133 846 41.9 8.ja 7 136 847 43.4 8.33 9 48.7 735 15.1 8.8b 70 48.7 735 13.9 9.0b 12 48.7 742 15.0 8.8b 13 48.7 74 1 14.7 8.9* 14 48.7 742 15.1 8.8b a Meazured during analysis of sample. Measured in reactor discharge line.
1-1/2" PIPE
Hypochlorite was discharged from the reactor through
a rubber hose to a drain. Provisions were made for a p H measuring point and a sample take-off. Analytical. Available chlorine concentration was determined by the thiosulfate method. Soda ash feed concentration was determined by titrating to a potentiometric end point with standard acid. Excess carbonate was determined by destroying the hypochlorite with neutral peroxide and titrating to the potentiometric end point with standard acid.
DRAFT T U B E
CHLORINE
Experimental Results
Redox Potential. T h e correlation developed between redox potential and available chlorine concentration is shown in Figure 3. T h e data were reproducible, and control end points for particular concentrations could be predicted by using the curve. At the reaction end point, the curve shows a slight inflection which is not as sharp as the inflection noted with caustic soda-based solutions. Figure 3 is a smoothed curve constructed from redox curves for 48.7, 117.3, 131, and 152-gram-per-liter sodium carbonate solutions. T h e portions of the curves beyond the inflection points have been eliminated. Since chlorinating beyond the inflection point would result in overchlorination and rapid decomposition, operating conditions for a particular available chlorine concen-
Table IV.
tration would be determined from the redox potential of Figure 3 and the raw material requirements of 'Table I. Redox curves for caustic soda-based solutions are generally given by redox potential as a function of excess hydroxide concentration ( 3 , 70, 72). In the work mith soda ash, no correlation was found between potential and excess carbonate Concentration. The degree of completion of Equation 2 Irould affect this relation. Capacity. The reactor operated satisfactorily over a range of flow rates from 20 to 130 gallons per hour for all concentrations tested. Process control was good throughout the range. Capacity appeared to be limited by the discharge line capacity.
Stability of Hypochlorite Solutions Was Studied as a Function of Time, Concentration, and Temperature
Initial Theoretical Available Chlorine, G./L. Initial Sample Na2C03, ATaHC03 7-Hr. Loss, 4-Hr. Loss, Temperature, No. G /L. G./L a Initial later % later 70 O 4* 131 94.8 40.0 35.8 10.5 28.4 29.0 23 10-lb 133.6 97.9 41.3 34.6 16.2 25.5 38.2 26 10-2c 133.6 98.8 41 . 7 36.8 11.7 33.5 19.6 26 14s 48.7 35.8 15.1 15.0 0.7 14.8 2.0 23 lld 133 98.1 41.4 35.3 14.8 ... 26 12-ld 136 102.9 43 4 24.2 44.5 .., 37 12-2d 136 98.4 41.5 26.8 35.4 *.. 37 a Assumes Equation 2 goes to completion. Maintained at initial temperature. c Chilled slowly, 7-hr. temperature 79' C., 4-hr. tpmperature 5' C. Cooled slowly to room temperature.
c.
...
... ...
~
VOL 1
NO. 2
~~
~~
APRIL 1 9 6 2
151
AVAILABLE CHLORINE, GPL. Figure 3. This curve shows the relationship developed between redox potential and the available chlorine concentration of hypochlorite solutions and can b e used to predict control set points
Reactor Configuration. During several runs, samples were taken from the electrode position and compared with samples taken from the discharge. These samples showed the reaction was complete a t the electrode position. At low flow rates, the electrodes could be moved closer to the chlorine orifice without affecting operation. At high flow rates, the electrodes could not be lowered because turbulence and bubbles caused erratic operation. The top section of the reactor was not required for successful operation. Pressure Operation. T h e reactor was operated under varying pressures up to 30 p.s.i.g. at the top of the reactor with no effect on operation. T h e total pressure was limited by the reactor working pressure and the feed pump capacity. Chlorine Losses. The vent gas flow could be passed through a caustic soda absorber to check for chlorine losses. Loss was essentially zero when operating under automatic control. Large bubbles passed through the reactor; however, this was due to carbon dioxide formed in Equation 1 rather than to unreacted chlorine. Chlorine Control. The slope of the redox curve in Figure 3 is such that automatic control is easily maintained. T h e data in Tables I1 and I11 show that uniformity of product is excellent and that reproducible results can be obtained. Hypochlorite Stability. Stability of 1 and 4 hours at various temperatures and concentrations is shown in Table IV. As expected, high strength solution was very unstable owing to relatively high bicarbonate Concentration. The bicarbonate concentration in high strength solutions was in excess of its solubility, and crystallization occurred upon standing. The half life of 40-gram-per-liter solution at room temperature of 22" C. was 8 to 14 hours. By slowly chilling the solution 152
l & E C PROCESS D E S I G N A N D DEVELOPMENT
over a 5-hour period to 3' C., the half-life was extended to 7 days. The temperature effect is also shown by the difference in 1-hour stability at 26' C. and 37' C. For any given soda ash concentration, the stability of the hypochlorite decreases as available chlorine concentration increases. Optimum stability is dependent on the end usage of the hypochlorite and how soon it will be used. Sodium Bicarbonate Formation. Because of low bicarbonate solubility, in hypochlorite solutions above 20-gram-perliter available chlorine, the bicarbonate would be expected to crystallize out if Equation 2 proceeded nearly to completion. Crystallization did occur in the discharge line and in sample bottles during 40-gram-per-liter runs, none occurred in 15gram-per-liter runs. No crystallization occurred in the reactor except when attempting to make higher than 40-gram-per-liter solutions. Complete plugging of the discharge line and partial plugging of the reactor occurred under this condition. T h e plugging could be cleared by flushing with water. Plugging should be no problem if the concentrations in Table I are used and if smooth pipe lines with no low spots are installed. Nomenclature
E = Potential E" = Standard potential F = The Faraday constant, 96,500 coulombs R = The universal gas constant, 8.314 joules
' K./gram mole T = Absolute temperature, ' K. n = Number of electrons involved in the ionic change of state under consideration a = Activity of a particular reaction component = Electron
Literature Cited
( 7 ) Matthews. J. M.. “Bleachine and Related Processes.” DD. 341-53, Chemical Catalog Co., Eew York, 1921. (8) Morton, D. S., Am. Dyestuf Reptr. 47, 372-4 (June 2,1958). (’9) ~, Nikolskv. B. P.. Flis. I. E., J . Gen. Chem. U.S.S.R. (Eng. . Transl.) 22, 1343-9 (1952). (10) Pye, D. J., J . Electrochem. SOC.97, 245-48 (August 1950). (11) Soci6t6 d’Electro-Chimie d’Electro-metallurgie et des Acieries Electriques d’Ugine, Brit. Patent 795,410 (May 21, 1958). (12) FYyandotte Chemicals Corp., IYyandotte, Mich., Michigan Alkali Tech. Service Rept. 20, September 1957. (13) Yorston, H. F., Pulp Paper Can. 33, 74 (1932). \
(1) Chem. Eng. 61, 128-30 (March 1954). (2) Duncan, E. P., Pulp &Paper 30,124 (Feb. 1956). (3) Fisher. H. S., Carlson, R. E., Western International Meeting of the Pulp and Paper Industry, Victoria, British Columbia, May 21, 1955. (4) Glasstone, S., “An Introduction to Electrochemistry,” p. 268, Van Nostrand: Princeton, N. J., 1942. (5) Hodgeman, C. D., et ai., “Handbook of Chemistry and Physics,” 34th ed., p. 1548, Chemical Rubber Publishing Co., Cleveland, Ohio, 1952. (6) Lister, M. IY., Can. J . Chem. 34, 465-78 (1956).
I
I
’
RECEIVED for review December 1 1960 ACCEPTED July 13. 1961 I
2,5=DIMETHYLPIPERAZINE SYNTHESIS FROM ISOPROPANOLAMINE W. K. LANGDON, W . W. LEVIS, JR., AND D . R . JACKSON Wyandotte Chemicals Corp., Wyandotte, Mich.
In the presence of hydrogenation-dehydrogenation catalysts, isopropanolamine undergoes a bimolecular cycloamination reaction to form 2,5-dimethylpyrazine, frans-2,5-di-methyIpiperazinet and cis-2,5-dimethylpiperazine. A study of reaction variables in liquid phase with Raney nickel as catalyst resulted in a practical process for making trans-2,5-dimethylpiperazine. The best yields of 2,5-dimethylpiperazine and the highest trans to cis ratios were obtained a t 1 2 0 0 p.s.i.g., 2 2 0 ” C., with 2.5 grams of Raney catalyst per mole and with a reaction time of 4 to 8 hours. Under these conditions isopropanolamine was almost completely consumed and the yield of mixed dimethylpiperazine isomers was about 75y0. The trans isomer content of this product was about 80%.
has been shown recently in the use of piperazines as chemical intermediates. No doubt this has resulted from the bifunctionality conferred by the two secondary amine groupings in these cyclic structures as well as from the development of new processes for making piperazine and carbon substituted piperazines from readily available raw materials such as ammonia, alkylenediamines, and olefin oxides. trans-2,5-Dimethylpiperazine, a white crystalline solid melting a t 118’ C.. has shown particular promise as a n intermediate for linear polyamides because of its symmetry. The cis isomer, a liquid freezing a t about 18” C., probably has limited use in polymers but is expected to be useful as a pharmaceutical intermediate. The development of a practical process for making 2,5-dimethylpiperazines from isopropanolamine (1amino-2-propanol), and particularly, the development of a method for controlling reaction conditions so as to obtain trans-2,5-dimethylpiperazine as the major product. is described. I n the presence of hydrogenation-dehydrogenation catalysts. isopropanolamine undergoes a bimolecular cycloamination reaction to form 2,5-dimethylpiperazine: ONSIDERABLE INTEREST
”’ H
Raney nickel was found to be an excellent catalyst for this reaction and was used in the study reported in this article. Previously, Bain and Pollard ( 7 ) obtained trans-2,5-dimethylpiperazine by heating isopropanolamine a t 250” to 275” C. in the liquid phase with copper chromite as catalyst and dioxane as solvent. Subsequently in these laboratories, reaction of isopropanolamine in the vapor phase over copper chromite was found to give 2,5-dimethylpiperazine, but the main product under these conditions was 2,5-dimethylpyrazine (3). More recently, Japanese workers have been active in this field (2, 4 ) . Although the net result is cyclodeh) dration, the actual course of the reaction is undoubtedly much more complex. Figure 1 illustrates one course that the reaction may take. In the presence of a hydrogenation-dehydrogenation catalyst, the first step probably involves dehydrogenation of isopropanolamine to form aminoacetone. Two molecules of this amino ketone would then cyclize to form a di-Schiff base. At atmospheric pressure, or at a few atmospheres of hydrogen pressure, the favored reaction is further dehydrogenation to form dimethylpyrazine. At higher hydrogen pressures, the favored reaction is hydrogenation of the di-Schiff base to form dimethylpiperazine. The mechanism shown is probably a n oversimplification since there are a number of ways by which amino acetone can react with itself as with isopropanolamine, and by which intermediates can be dehydrogenated or rehydrogenated. Although pressure is a n important variable, reaction time, reaction temperature, and amount of catalyst also strongly VOL. 1
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