Chloramine Equilibria and the Kinetics of Disproportionation in

16 C h l o r a m i n e E q u i l i b r i a a n d the K i n e t i c s of D i s p r o p o r t i o n a t i o n in A q u e o u s S o l u t i o n 1

2

Downloaded by UNIV OF SYDNEY on August 14, 2015 | http://pubs.acs.org Publication Date: January 12, 1979 | doi: 10.1021/bk-1978-0082.ch016

EDWARD T. GRAY, JR., DALE W. MARGERUM, and RONALD P. HUFFMAN Department of Chemistry, Purdue University, West Lafayette, IN 47907

Introduction The chlorination of water as a means of disinfection contin ues to enjoy popularity because of the low cost, simplicity, and great effectiveness of the technique. Although the process of disinfection is not well understood, it is known that as chlorine is introduced into waste water, the amines present react rapidly with the chlorine to form chloramines. Chloramines have been shown to be toxic to fish when present at lower than ppb levels (1). However, quantitative analytical procedures sensitive at this level are not available. The high toxicity to aquatic life has prompted investigations in this laboratory aimed at extend ing existing analytical techniques to lower levels of sensitivity. One question which must immediately be addressed in this work is the effect of changing pH upon the distribution, or the very existence, of the chloramines present in a sample taken for analysis. In the case of ammonia, for example, the equilibria of the chloramine species can be expressed by eq 1-3. Equation 1 Κ NH + HOC! ( 3

+

H Y D

> • NHC1 + H0 2

2NHC1 + H JVj?_> 2

(1)

2

NHC1 + NH 2

+

4

(2)

'Current address: Dept. of Chemistry, U. of Hartford, W. Hartford, CT 06117 Current address: Upjohn Co., P.O.B. 685, Laporte, TX 77571

2

0-8412-0461-6/78/47-082-264$05.00/0 © 1978 American Chemical Society In Organometals and Organometalloids; Brinckman, F., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1979.

16.

GRAY

E T

3NHC1 + H 2

Chloramine Equilibria

AL.

+

>

X p T

c

2NC1 + N H 3

might also be expressed, NH

3

+ OCl"

(

» M

Values o f Κ

Η γ ο

(3)

+

i n basic media, as shown i n eq 4.

NH C1 + OH"

(4)

2

,

K^p, and Kpj have h i s t o r i c a l l y been

c u l t to o b t a i n . C o r b e t t , Downloaded by UNIV OF SYDNEY on August 14, 2015 | http://pubs.acs.org Publication Date: January 12, 1979 | doi: 10.1021/bk-1978-0082.ch016

4

265

M e t c a l f , and Soper (2j

diffi

reported a value

for K o f 3.6 χ 1 0 M" (15°C, Κ = 4.7 χ 1 0 ' Μ , Κ of πτυ Q w a H0C1 = 2.7 χ 10 M) while Anbar and Yagil (3) o f f e r evidence 9

y v n

1

1 4

2

that v f the number can be measured at e q u i l i b r i u m , the value greater than 1 0

1 5

M" -8

for H0C1 = 4 χ 10

1

(27.3°C,

M).

M"

μ = 1.0 M, c a l c u l a t e d using a Κ

Bridging t h i s large gap, Granstrom

c a l c u l a t e d the value for K

is

(4)

from k i n e t i c data to be 2.4 χ 1 0 ^

u v n

(25°C, μ = 0.45 M, K o f H0C1 = 2.9 χ 10~° M). a The e q u i l i b r i a shown i n eq 2 and 3 have not been evaluated

experimentally because o f the reported l a c k o f s t a b i l i t y o f the p o l y c h l o r i n a t e d species

(2_,

able to estimate values o f 1 0

6_, 7). 6

M"

1

However, J o l l y (8)

and 1 0

rium constants K^p and Kp^, r e s p e c t i v e l y . data o f Chapin (5)

4

M"

1

was

for the e q u i l i b

This was done using

and C o r b e t t , M e t c a l f , and Soper (2)

even

though the l a t t e r authors conclude that t h e i r data do not r e p r e sent an e q u i l i b r i u m c o n d i t i o n . (5J,

Morris (9)

Using the same data o f Chapin

c a l c u l a t e d K^p to be 2.3 χ 1 0

7

M' . 1

Using the

h y d r o l y s i s constant o f C o r b e t t , M e t c a l f , and Soper {2), estimates o f l^p and K p , J o l l y (8) T

electrochemical

and his

was a l s o able to estimate

p o t e n t i a l s for the ammoniacal chloramines.

In the present work, a l l

chloramine s o l u t i o n s are

prepared

with excess amine present and are manipulated with the use o f a double-two-jet in stopped-flow

tangential mixing system, s i m i l a r to those used spectrophotometers.

This assures

rapid and

e f f i c i e n t achievement o f homogeneity ( l e s s than 100

\is).

Chloramine s o l u t i o n s produced in t h i s manner are s i g n i f i c a n t l y

In Organometals and Organometalloids; Brinckman, F., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1979.

ORGANOMETALS AND ORGANOMETALLOIDS

266 more s t a b l e

than those produced by hand m i x i n g ; they have e x h i

bited no l i g h t s e n s i t i v i t y ; and experimental consistently duplicated.

r e s u l t s can be

Monochloramine s o l u t i o n s prepared

t h i s manner are s t a b l e at room temperature,

undergo

in

dispropor-

t i o n a t i o n to NHC1 without apparent side reactions at pH NHC1 + N H

S

2

4

(10)

+

The rate expression f o r t h i s microscopic r e a c t i o n i s d [ R N C 1


d

2

]

+ = k

2

t

D I S

drRNHPll

[ N H C l ] [ N H C l ] = l/2 ( 2

3

+

d

L g"J)

(11)

R

In terms o f [NH C1] (=[NH C1] + [NH C1 ]) 2

T

2

3

+

(12) -d[RNHCl] d

t

where

s

2

III.

K

[ HH + ] [ N +H

NH C1

+

Values o f 2

D(ΊI S+ HK [ H ] 2)

2C 1 ]

T

i s the protonation constant for monochloramine (eq 1 3 ) .

NH C1 + H

NH ,

2 k

3

(13)

+

and kpj^ were resolved f o r chloramines o f N H , C H 3

3

β - a l a n i n e , g l y c i n e , and g l y c y l g l y c i n e as compared i n Table The values

for g l y c y l g l y c i n e are the r e s u l t o f a c o n s i d e r

able e x t r a p o l a t i o n but are c e r t a i n l y as accurate as they would be if

the 0.5 M i o p i c strength r e s t r i c t i o n had been broken i n order

to achieve the 2-5 M a c i d necessary to obtain data i n the region of

the k i n e t i c maximum. It should be noted that these values o f k p

small and the

I S

are rather

values are higher than expected (12).

protonated monochloramines w i l l

Hence

e x i s t at low pH f o r a s i g n i f i c a n t

time period at low t o t a l concentrations o f RNHC1.

Accordingly,

e q u i l i b r i u m constants and electrochemical p o t e n t i a l s are included for the protonated species i n Tables III.

I and I I .

D i s t r i b u t i o n o f Ammoniacal Chloramines Figure 6 d i s p l a y s the c a l c u l a t e d d i s t r i b u t i o n o f m i l l i m o l a r

c h l o r i n e i n a f i v e - f o l d excess o f ammonia as a function o f pH. This i s analogous to the c o n d i t i o n s used for the preparation o f stock monochloramine i n t h i s work.

Besides the complete forma

t i o n o f monochloramine above pH 9 . 5 , i t i s i n t e r e s t i n g to note


274

ORGANOMETALS

AND

ORGANOMETALLOIDS

Figure 4. Approach of a mixture of NH$, β-alanine and their monochloramines toward equilibrium. (1) initial spectrum; (2) equilibrium spec trum, [HOCl] = 1.31 X 10 M, [NH ] = 0.059M, [β-aknineL· = 0.0022M, -log[H+] = 9.08,1-cm path length, μ = 0.50M (NaClO ) 25.0°C; (3)... fully formed N-chloroβ-alanine.


S

S

T

h

Figure 5. Dependence of k (M' sec' ) on [H*] for the disproportiona/—d[RNHCl] .2k [RNHCl] ^ tion I dt of monochloramine (Φ), N-chloro-βalanine (A), and N-chloroglycylglycine (O). ([Chloramine] — 1.06 X JO M, μ - 0.50M (NaClOt), 25.0°C).

230

270

310

Wavelength, nm

1

oba

1

ob8

T

s

Figure 6. Distribution of chlor amines in aqueous solution ([NH ] — 10 M, [HOCl]τ — 2.5 X JO M, O.50M (NaCl0 ), 25.0°C.) The decomposition of NHCl, above pH 6is neglected. S

-log

[Hi

2

T

s

4

-tog[H+]



16.

GRAY

E T AL.


275

Figure 7. Distributions of chloramines in aqueous solutions with [NHS]T — [HOCl] , 25.0°C, 0.50M (NaClOJ. The decomposi tion of NHCl, above pH 6 is neglected: (a) I0" M, (b) 10' M (c) 10- M. T

3

6

9

9

Figure 8. Distributions of chloramines in aqueous solution with a constant ammonia concentration and varying [HOCl] ([NH,] — ΙΟ", 25.0°C, μ — O.50M NaClOJ. T

6

T

The decomposition of NHCl above pH 6 is neglected: (a) [HOCl] = iO M, (b) [HOCl] = I0' M, (c) [HOCl] = I0- M, (d) [HOCl] = 10 M. g e

T

7

logCHl

T

8

T

T

9


276

ORGANOMETALS AND ORGANOMETALLOIDS

Table

III

Protonation Constants o f Monochloramines and the Rate Constants for D i s p r o p o r t i o n a t e -d[monochloramine]/dt (μ

= k [monochloramine][H-monochloramine DIS

]

= 0.50 M (NaCIO,), 25.0°C)

Amine


from

K , M" R

k

1

p i s

,

M"

1

Methyl ami ne

36

60

Ammonia

28

980

3-alanine

2.4

Glycine

0.37

1200

Glycylglycine

0.21

1100

sec

330

that the e q u i l i b r i a do not allow f u l l

formation o f dichloramine

or t r i c h l o r a m i n e under these c o n d i t i o n s .

Figure 7 i l l u s t r a t e s

the c a l c u l a t e d d i s t r i b u t i o n s o f chloramines at e q u i l i b r i u m mixing equimolar hypochlorous a c i d and ammonia at concentration l e v e l s .

after

various

These three plots present the changes

which occur upon d i l u t i o n o f chloramines under these concentra t i o n and pH c o n d i t i o n s .

The t r a n s f e r o f c h l o r i n e from nitrogen

to oxygen ( h y d r o l y s i s ) i s the predominant o v e r a l l

process.

Another facet o f the chloramine e q u i l i b r i a i s presented i n Figure 8.

In t h i s figure

the t o t a l ammonia concentration i s held

constant at 10"^ M and the a v a i l a b l e -9 to 10

M.

As the a v a i l a b l e

c h l o r i n e i s varied from 10"^

c h l o r i n e concentration d e c l i n e s ,

c h l o r i n e s i n g l y bound to oxygen or nitrogen becomes preferable to p o l y c h l o r i n a t e d nitrogen m o i e t i e s .

Above pH 6 NHCl2 i s not s t a b l e

r e l a t i v e to s e l f o x i d a t i o n - r e d u c t i o n and therefore

these

distri

butions represent only the maximum p o s s i b l e values o f d i c h l o r amine. Acknowledgment. National

This i n v e s t i g a t i o n was supported by

Science Foundation Grant CHE75-15500.


16.

gray et

al.


277

References and Notes


1. 2.

Merkins, J. C., Water Waste Treatment J., (1958), 7, 150. C o r b e t t , R. E., M e t c a l f , W. S., and Soper, F. G., J. Chem. S o c . , (1953), 1927. 3. Anbar, M. and Yagil, G., J. Am. Chem. S o c . , (1962), 84, 1790. 4. Granstrom, M. L., Ph.D. T h e s i s , Harvard U n i v e r s i t y , 1954. 5. Chapin, R. M . , J. Am. Chem. S o c . , (1929), 5 1 , 2113. 6. M o r r i s , J. C., "Kinetics of Reactions Between Aqueous Chlorine and Nitrogen Compounds" in "Principles and Applications of Water Chemistry," Faust, S . D. and Hunter, J. D . , e d s . , W i l e y , Ν. Υ., 1964, pp 23-53. 7. M e t c a l f , W. S., J. Chem. S o c . (1942), 148. 8. J o l l y , W. L., J. Phys. Chem., (1956), 60, 507. 9. W e i l , I. and M o r r i s , J. C., 116th National Meeting of the American Chemical Society, Atlantic C i t y , N. J., September, 1949. 10. G a l a l - G o r c h e v , H. and M o r r i s , J. C., Inorg. Chem., (1965), 4 , 899. 11. L a t i m e r , W. Μ., "The Oxidation States o f the Elements and T h e i r Potentials in Aqueous Solutions," 2nd e d . , Prentice H a l l , I n c . , New York, 1952, pp 53-55. 12. W e i l , I. and M o r r i s , J. C., J. Am. Chem. S o c . , (1949), 71, 3123. Received August 22, 1978.


Chloramine Equilibria and the Kinetics of Disproportionation in

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