SEPTEMBER. 1955
ERVIN COLTON Georgia Institute of Technology, Atlanta, Georgia
MARK M. JONES University of Illinois, Urbana, Illinois
E A R L Y in this century Raschig ( 1 ) found that when dilute solutions containing equimolar amounts of ammonia and sodium hypochlorite are mixed, the resulting solution does not give the characteristic reactions of the hypochlorite ion as would be expected if ammonium hypochlorite were formed. He also noted that the evolution of nitrogen, from the oxidation of the ammonia, is very slow. On the basis of these observations Raschig postulated the presence of monochloramine, NH,Cl, in these solutions, and showed subsequently by analysis that the active oxidizing agent contains nitrogen and chlorine in the ratio of 1/1. Although he never succeeded in isolating pure monochloramine, Raschig studied many of the reactions of this elusive substance using aqueous solutions prepared from ammonia and sodium hypochlorite. In subsequent investigations Raschig discovered that monochloramine can react with ammonia in basic solutions to form hydrazine in accordance with the equation:
NHGI
+ NH3 + OH-
-
+ H 2 0 + C1-
N2H4
This method of synthesis is still the most convenient for the preparation of aqueous solutions of monochloramine ($9). Other reactions in which monochloramine is formed include the following: ( a ) The reaction between tertiary hutyl hypochlorite and ammonia (3): (CH3)GOCI
+ NHa
-
NHKl
+ (CH2)&OH
( b ) The reaction between chlorine and ammonia, either dissolved in water (4, 5) or in the gaseous state with nitrogen as a dilnent (6, 7) : 2iYH1
+ CI?
-
NHICI
+ KH,CI
( c ) The hydrolysis of dichlorourea (8):
( d ) The reaction of nitryl chloride and ammonia at -75'C., in which monochloramine is reported to be one of the products (9) :
-
NOzCI + 2SH1 NHzCl + NHciY02 Hydrazine is now prepared commercially by a process which essentially embodies the original observations of (e) The acid hydrolysis of potassium chloraminoRaschig-a process usually called the Raschig synthesis. sulfonate (10): Because much attention recently has been focused on NHCl SO,K + H 2 0 N H D + RHSO, hydrazine as a specialty fuel, monochloramine, the key intermediate in the Raschig synthesis of hydrazine, Reactions ( a ) , (b), and (c) have been used to prepare assumes an important role commercially. Further- monochloramine as an intermediate for the preparation more, the bactericidal properties of monochloramine of hydrazine. Reactions (d) and (e) are of academic make this ammonia derivative one of the most widely interest. Because of the instability of monochloramine, used chemicals in the purification of water. For these i t is generally employed in aqueous solution or in ot,her reasons it seems appropriate to summarize certain as- suitable nonaqueous solvents such as diethyl ether. pects of the chemistry of monochloramine.
-
PROPERTIES
PREPARATION
Monochloramine is formed. as was shown bv Raschie. -, through the interaction of ammonia and inorganic hypochlorites in accordance with the equation:
Monochloramine has been isolated only once (11) by uassinc the vauors of an aaueous solution over uotassium carbonate and condensing the product a t liquidair temperature. The pure substance forms colorless
.
-
JOURNAL OF CHEMICAL EDUCATION
488
hypochlorite ion. These authors (18) also confirmed the earlier workof Chapin (19) who showed that theproducts of the reaction between ammonia and hypochlorite depend upon the pH of the solution. Below a pH of 3, nitrogen trichloride is formed; between a pH of 3 and 5, dichloramine is obtained; and above a pH of 8, monochloramine is produced. These represent, of course, the principal products formed in each pH region. There is an equilibrium between mono- and dichloramine in diethyl ether > isopropyl ether > B,B'-dichlorodiethyl ether > benzyl ether > sym.-tetrachloroethane > butyl ether > ohloro- the range of pH values of from 5 to 8 and a similar form > benzene > chlarobenzene > carbon disulfide > carbon equilibrium between nitrogen trichloride and dichlortetrachloride > low-bailing petroleum ether > eyclohexane amine near a pH of 3. crystals melting a t -66°C. and decomposing a t slightly higher temperatures. Monochloramine is soluble in water and in diethyl ether, the distribution ratio between the two solvents being approximately one (11). Qualitative studies on the extractability of monochloramine from aqueous solutions with various nonaqueous solvents show the following order of decreasing effectiveness (12) :
+
+
N&CI O H e NHI OCIMonochloramine is highly useful in water treatment 2NH1Cl H + a NHCI* + NH4+ as a means of destroying pathogenic bacteria. Ingols and co-workers (IS) have shown recently that, in comparison with hypochlorous acid, monochloramine reMonochloramine solutions are difficult to preserve quires much more time and larger concentrations of for extended periods, the decomposition products being oxidation capacity to bring bacterial death. dependent upon the pH of the solution. Monochloramine may be looked upon as an ammono hypochlorous acid in which the hydroxyl group is re- ANALYSIS placed by its nitrogen equivalent, the amide group (14). Three common methods are usually employed for the The pK for monochloramine is estimated (15) to he analysis of monochloramine solutions: (1) iodimetry, 15 2 in comparison with a pK of 8 for hypochlorous ( 2 ) colorimetry, and (3) spectrophotometry. The first Hence, only in strongly alkaline solutions would acid. monochloramine be expected to be appreciably ionized. two procedures are based upon the ox'dizing character of monochloramine, while the latter method depends NHsCI H i NHClupon the physical constitution of the molecule. The first method of analysis uses the reaction between The oxidation potentials of monochloramine have monochloramine and iodide ion in acid solution to generbeen calculated (15) : ate iodine, which is then titrated with thiosulfate. The C1- + NHlt a NH&I + 2Ht + 2e-; ED = -1.45 v. reaction proceeds according to the equation: C1- + OH- + NH. ~5NHCI H.0 + 2ec; EO = NH,Cl + 2Ht 21NH4+ + b + C1-0.75 v.
+
*
+
+
+
-
Best results are obtained if the monochloramine solution is introduced below the surface of the acidified iodide solution, to decrease loss of iodine by volatilization. The second method for the analysis of monochloramine is the so-called Palin test (20) in which an instantaneous red color develops when monochloramine, in a phosphate buffer of pH 6.8, is mixed with iodide ion and p-aminodimethylaniline. The color fades rapidly, and standard conditions must be observed for consistent results. I n a more recent method Palin (21) employs a solution of neutral o-tolidine. A blue color apHOCl NHI NH&l + H 2 0 pears in the presence of monochloramine after the addiThese authors (17) also studied the base strength of tion of iodide ion, the blue color then being titrated with monochloramine and propose an equilibrium constant ferrous ammonium sulfate. of 1 X for the reaction: The third procedure for the determination of monochloramine is dependent upon the characteristic abNH,CI + H 2 0 a NH3CI+ + OHsorption band of this substance a t 2430 A., E = 458. By comparison, the value for ammonia is 1.8 X 10-5. Recent studies (22) reveal that over the pH range 9-11 A reaction more intimately connected withtheRaschig monochloramine obeys Beer's law and can be determined synthesis is the equilibrium: in concentrations as low as 10-4 M. NH&l + OH- e NH3 + OCIREACTIONS A number of the reactions of monochloramine have Corbett, Metcalf, and Soper (18) used a spectrophotometric method to determine the equilibrium constant been investigated in some detail. One of the most imfo? this reaction and arrived a t a value of 1.6 X 10-3. portant reactions, from a commercial point of view, is This means that only in very strongly basic solutions that between monochloramine and ammonia to form does monochloramine decompose into ammonia and hydrazine (see equation in introduction). I t has The potential for acid solutions is in agreement with the fact that monochloramine is a stronger oxidizing agent than bromine (for Br--En, E L -1.09 v.). The potential for alkaline solutions is in agreement with the fact that monochloramine is a weaker oxidizing agent than hypochlorite (for C1:OCl-, E i = -0.89 v.). The kinetics of formation of monochloramine from ammonia and hypochlorous acid have been investigated by Weil and Morris (16) who find the reaction to be second order in accordance with the equation:
+
-
SEPTEMBER, 1955
been found, however, that hydrazine and monochlorarnine react according to the equation: 2NHzCI
+ N,H4
-
NB
+ 2NHdCI
I n order to realize substantial yields of hydrazine from the ammonia-monochloramine reaction it is necessary to decrease this yield-reducing reaction between the unreacted monochloramine and the hydrazine that is formed. Since traces of heavy-metal cations, especially copper, have been found to have a marked ratalytic effect on the monochloramine-hydrwine reaction (23), the addition of gelatin to the synthesis solutions serves to chelate the undesirable cations and thus permits excellent yields of hydrazine. Sisler and co-workers (24, 26) have investigated the monochloramine-ammonia reaction in liquid ammonia and in pure water. NHKl
+ 2NH8
-
NYH,
+ NH&I
I n the former solvent at 10O0,yields of hydrazine in excess of 80 per cent of theoretical (based on monochloramine) are obtained for a mole ratio NH,/NHsC1 of approximately 400. Using pure water as the solvent these workers have shown that monochloramine and ammonia yield hydrazine in excess of 80 per cent of theoretical, but that the over-all yield is dependent upon the monochloramine concentration. Raschig (26) reported that phenylhydrazine results when aniline and monochlorarnine react:
+
C6H6NHn NH&I
-
CsHsNH-NH*
+ HCI
Audrieth and Diamond (97) have demonstrated that monochloramine will react with various primary alkyl amines in basic solution to form the corresponding Nsubstituted hydrazines, isolated as hydrogen sulfates. RNHI
+ NH&L
-
+ HCI
RNH-NH?
I t is interesting to note that, in contrast, monochloramine reacts with certain amino acids a t pH 8 to form the corresponding N-chloro derivatives (IS). This reaction has been observed with alanine, glycylglycylglycine, and tyrosine, the products being identified spectrophotometrically. RNH1
+ NHnCl
-
RNHCI
+ NHr
McCoy (28) has presented evidence for the presence of hydroxylamine as an intermediate in the decomposition of monochloramine by hydroxide ion. NHnCl
+ OH--
NHzOH
+ C1-
The presence of hydroxylamine was demonstrated by the isolation of cyclohexanone oxime. Monochloramine reacts with certain aldehydes to form solid organic derivatives, ald-chlorimines, of the general formula RCH=NCl (29). RCHO
+ NHzCI
-
RCH=NC1
+ HsO
These crystalline compounds of well defined melting points may he used to establish the presence of macro amounts of monochloramine.
SUMMARY
Some of the physical and chemical properties of monochloramine have been presented. The usefulness of this highly reactive molecule in water purification and in hydrazine synthesis serves to illustrate that monochloramine can no longer be looked upon as a laboratory curiosity. ACKNOWLEDGMENT
The authors are grateful to Professors L. F. Audrieth and R. S. Ingols for generous discussions and helpful suggestions. LITERATURE CITED (1) RAscHro, F., "Schwefel- und Stiokstoffstudien," Verlag Chemie, GMBH, Berlin, 1924, pp. 5IF78. This book is a summary of 30 years of Rsschig's researches. (2) BOOTH,H. S., "Inorganic Syntheses," McGraw-Hill Book Co., h e . , New York, 1939, Val. I, p. 59. (3) AUDRIETH, L. F.,E. COLTON,AND M. M. JONES,J . Am. Chem. Soc., 76,1428 (1954). (41 SCHONBEIN. C.. J. vmkt. Chem.., 84. 385 (1861). , i5j Cnoss, C. F., E.J.'BERAN,A N D J. F. BRIGGS,J . SOL Chem. Ind., 27,260(1908). (6) MATTAIR,R., A N D 1%.H. SISLER,J. Am. Chem. Soc., 73,1619 11961) --- - , (7) SISLER,H. H., F. J. NETH,R. 8. DRAGO,AND D. YANEY, ibid., 76,3906 (1951). (8) DATTA,R. L., J. Chem. Soe., 101,167 (1912). (9) BATEY,H. H., AND H. H. SISLER,J. Am. Chem. Soe., 74,3408 (1952). (10) FRIEND,J. N., "Textbook of Inorganic Chemistry," Charles Griffin and Co., Ltd., London, 1928, Vol. 6, p. 117. (11) MARCKWALD, W., A N D M. WILLE,h r . , 56, 1319 (1923). Unpublished results oh(12) COLTON,E., A N D L. F. AUDRIETH, tained on the Hydraaine Program, University of Illinois,
.
\
~~~
.
.
lo
