ori

n.et;.ics

A

to

of iGd

of

fOI' the;: of

Doctor of Ph os

;ject

Harva:r·d vers:i

Co:pyr t rHSElrvod by

PREFACE

of i.nt reac on

ions has been a

zle fj.eld of water treatment, :cesea,:rch

s thes ents EL cox1c:entrated to

eluc ts (;f t s of tho

action a:nd to some o:f its

was conductt.:d

at versity f a.nd was c;d by a

ic Heal

'fhe to d.e to

fes;::.or J" Mor.c 9 h or~ for the invaluabJ.. e

gu e dur the ent cou.rs e of res ~ and

of 1na:nuscr .]• +c' .. ,.;>.)

N, Butler, and on the

c ttt:e; to and Mr, John

Gaude-t tc for n.g the; c used

; and to my e-n_c

a.nd patient of mantiE1 pL

' l

~ IVla~::;;sachusetts

TABIJl~ OP CONTENTS.

IIIi3T OF FIGUHES ~

IIIST OF rr./tBLES G

SYNOPSIS •.•

CHltPTJ~H

CHAP'I'ER IL

CHJ~OHINA1.I'ION OF WNr.EH AND .EFFECrrs OF AMMONIA NIT'R.OGEN ••• e ••••• ~ ~ • ~ 0 • , •• ¢ •

ion of Wate:c

Waters

C ~ Format

D~ Br es ef th

DEVELOPMEN'T OF A MECHANIST

A~ Charc"J.(:;te:c:i s t of ex

Reaet Schernes

B. ly 1 ' J.\ ' M.

2 ' J ' H. Ross urn

3' A. ~r. LJ~ ~ J' c, Morr & I, VI, Weil

" \,._., ri> of w s Data D&

1 ' R.eaet Sc;h Bme ar1d S

2 ~ Sto of the Break:-po Roac

of Compu.te:e S 1at:ion

F ~ Pr Model

1~ Gen it 2, Results and F ngs

:zci

XJ .. l

1-l

1.. "'"1

1

1.

1-12

1-22

2

2

3

2-18

III

CHAP'l'ER IV,

A.

2.

c.

for

b, NOT-f''AS

c~ DPD-F'AS

sware

Water

b~ Chlo:r

c: ~ Amm o:rd. tnn

d.il} Bu

R

ffe:ce:n_ .. ti

tions

ions for DPD-·~I-?AS

!Vleasur(~ment

EXPEHI!\IJEN'I'Al, KINETIC S'l'UDIES

A~ Des of

11!1

Concentrat

tem

n of

on Method

ed

on.s

5. Molar Ra o of :.tne t~o

c. L t of

••

1

5 J-7

.3-10

10

1

2

2

3

2

2

CHAP'rEH v.

C !LAP'I'EE VI •

2 • Effect of pH • • . . • • • • . . . . . • • • . ' ' J, ,, '~ t of Ammonia

Cone . • . • . ' ' . " . ~~ . .Effect of Buffer System

F~EVISION AND EXPANSION OF ~~HE REAC~~riON IViODEI.

13,

c.

on of

l • e Kinetic Data and

2 ~

tions

.ion of k5 on the

l. Available Kinet

1 '

I,imi tat i:i.'•.';'+ .Ln.} V

Reac

of frorn

on tho

o:r1s and

Ki

on

2, Nitr·ot~en ide Format

on Data

(''

' ' . . L~--22

IJ·-23

1+-· 32

l

3

5-21.

3

COMPARISON OF' MEASURED AND COMPU'PED BREAKPOINT REAC~1 ION FEA1rUl~ES • ~....... • ~ ~ * •• 6

A~ Cu:erent 3

l * Sto

NCl3

J ~ T{inet

h, Max.

c,,

c o:ns

Formc.;.t of End

6

6 Cone ons

i:n Free Ghlo.r

0

q _l,J. on

l e Cor1dit

2~ Sto.i

b~ Max

of

in NHC of NC1")

J

and

erttrat

J ~ ~ 6~1.3

6mlLc

6 6

' . ' 6ml'l 17

CliAP'rEH VII, GENERAL EVALUATION , , , , •. , . , • , , ••. , • , • , , , • 'l

lt ~ ~'he Computer ' ' ' ' ' . ' ' ' . ' ' ' ' ' ' . ' ' ' B. The lVlaximum Cone ' ' ' ' " ~rh.n v' ' ' ' ' ' ' ' ' ' ' ' ' ' ' D, The !T1 e ,, ' ' ' ' ' ' ' ' ' ' . '

IOGRAPHY ~·. ~· .. ·~ ... ~,~·."." .. ~· ·~ .. ~~~· .. ~-" .... ~

'l '" i}

'l-5

API'ENDIX J. COMPTYC:E:H. PHOGRAl.VI FOR COMPUTATIONS OF

BEEAKPO IN'r KINE'J' ~~,.~~···~······~~···

APPENDIX II, ~lPECIFIC RA1l 1ES FOE COlVJPUTA1'IONS OF

BHEAKPO IN'r KINETICS •••• , •••.••..•. , .•• , A-9

APPENDIX III, IllfiENTAI, DAS'A ON ICS OF BHEAK-

POIWr REACTION BETWEEN AQUEOUS ClH~OlUNE

AND AMMONIA ••••••• t "~. ~ •••• ~ ~ •••••••••• A·.,lLV

vi

F

2-l

2-5

2

2

2

OF' FIGURES

Ti

Development of th.o th Chlor to Ammonia Molar Hatio to 1,62 and l,

vd.th in~":) to

vii

1. -12

1·~14-

to 1.62+0.04 ( led t ) ~o~·~~~ ;~,= .. 10

of the

of the Ammonia Nlolar Ha

6.5.

c 0/N0~1.6o

iminary Computat pH 7,5, C0/No~1.60

c:n. vs ~

tc: l . 82

til to 2,37

in~s at

•s Data at

1 1'11' s

f ' . t H 6- 7 °0°"

tc1

to

o i1..mmon1a a· p ~ w t:, vi! N w Co/N'o 1 ® B ~ G Ol M " ~ ~ ~ ~ a ·• .. ~ ~ ~ " ~ ~

Chlorination of Ammonia at pH 6,7, l5°C, No 1 mg/1 N, Co/No 1.8, Pt ,01 IV! •••••••••• I.'

Chlorj No 1

Ch1or No 1

No

of Ammo at pH 6~?w l N, Co/No 1.8, Pt .01 M

of ArnmonJ.a, at pH 6 ~ 7 ~ N, c0/N0 L8, Pt .01 l\1

0 5 c 1

ion of Ammonia at 7.0, 2 ·p·/J ,,, " /!! < Q 01 lVi mt.:>~ . l ~ ~ \_, o \l o -L @ ,_) w ~ "'· • J

on of Ammon at 7. 0, 1 mg/1 N, c 0/No 1.8, , ,Ol M ,,

of Ammonia at pH 7.0, l0°C, N, Co/N0 1,8 1 .01 M •m~•·~•••••e•

on of Ammo at pH '(,0, No 1 mg/1 N, Co/No 1.8, Pt ,01 lVI

2-11.

2-24

2-25

2-2?

2-28

4-?

L¥-1 0

2

Le-1 3 on of Ammonia at pH 7,2, 20°C,

N, Co/No 1.8, P.t .01 M ···~~•~m&~··· 4-14

Figure

3

6

7

1~.~ l 13

Lf··21

5

!1· ;,v, '\''''"

Chlor No 1

Chlor No 1

~'i.tle

of Ammon.ia at pH 7 ~ 2 ~ N, Co/N0 1,8, Pt .01 M

of Ammonia at N, Co/No 1,13, of Ammonia at pH 7,2,

N, Co/No 1.13, Pt .01 M 0 5 C,

t of at pH 7.0, noOn ~ \~ ~ No 1 I 5 mg/1 N' Co/No 1 . fl' Pc • 01 Nl ·~~~~~9~~"

Chl t:lon of Ammonia at pH 7 I 0, No 1.5 mg/1 N I Co/No 1,8' .01

Chl.or of .Ammonia at 7. 0 ~ No 1.5 N, C0/No 1,8, ,01

tion of Ammonia at pH 7,0, No 1.5 mg/1 N, Co/No 1.8, Pt ,01

5°C~ !V! ~··•~t""*~

Chlorinat No 0.25

in at No o~s

Effect of

E

Ch1or· No

Act;

f Ammonia at pH 6.7, 20°C, N ~ IN 1 8. ~ v Q/ - () ... ~ - ~

of Ammon at 6. 7, 20°C, N' Yl /j\j j Q D 0' JW ·W '"'01' () ~&()~ .i.'t .,.L VJ ~~0~~~~·~~

on Chlor ion of , Co/No 1.8, .01 M.

of Ammonia at pH 7,0, 20°C, N ~ Co/NO l ~ f:3 ~ P t ~ 01 lV1 ~ • ~ ~ g ~ • ~ ~ ~

Cone tion on ChJ.orinat of at pH 7.0, , co/No 1.8, .01 M.

of Ammonia at Nr c0/N0 lfflBf

7 ~2 ~ ;:~o 0 c ~ &O:L l'v1 ~@~~w~~~~~

on Chlod.nat of , c0/No 1.8. ,,01 Nl~

?n16 to 7~22~ o L7,

ii

I' age

5

L~.,.16

7

4-20

7

0

3

5

7

5--9

0

l

2

3

Lf

5

t: u

7

8

9

.5-20

2

/i.Ct for Forrnatior1 o:f 5-19

Compu 2ooc~

int c0/N0 1.8

1nt

Computation o:f Breakpoint tooc, No 1 mg/l N, c0/N0

L8

a.t pH 6 m 7 ~

s at pH 6 .• 7 ~

•1 ("' C' "" ·t pl'.r 0/ '7 .L . ...-k) {:;,t . .c '-- # 1 ~ 1,.8 ~~~~~$~$·~~ .. ~~~~9~

Int c0/N 0 1~8

, . ... ~~$~~~·~""~*~,~~ _)

Computat 15oc, No

C of

r:> at pH 7 . 0 ~

to°C, No 1 mg/1 N, Jnt

c0/N0 1,8 ("

~ ~ ~ • ~ $ 1! ~ ~ ~ ~ $ ~ ~ @ * ~ :J""'

c of Breakpoint s at 7.0. ~ No l mg/1 l\f ~ Co/No 1 m 8 ~ " * @ ~ ~ ~ ~ ~ m •• " ~ ~ • & ~

of Break~point at pH 7. 2. w ·f o·/-1 N·' r1 /!" 1 8 .~.- mo/ ..... L ~ ,..,.Ol ·~o ·"' ~ ... ~ ~ ~ ~ ~ " ~ ~ ~ ~ ~ « ~ ~ ~ ~ G

c

1

1. 5=36 C tahtior1 of

, No 1 Com·putat of Br

5°C ~ No l m.g/1 Computation of

20°C, No 1,.5

at pH 7,2,

at ?~2~ 0/N 0 1 * s & @ ~ ~ ~ ~ & • ~ • ~ •• 4 s • ~ ~ .5- J s

Kinet s at 6.7~ N, Co/No 1.8 . "••• ... ··~ ... 5-40

Computat of Breakpo Kirwt 20°C, No 0,.5 mg/1 N, Co/No 1,8

Computa.t:Lon of 2ooc ~ No 0 "~.5

lnt Ki:net N, Co/No 1,8

of int K e·tics at pH 7.2~ o " 5 mg/1 N w co/No l . 8 ~ . ~ ~ ~ ~ . * , ~ o & ~ ~ ~

c Nw Co/No 1~8

s at pH 7,t',

int Kinet vs , 's Data at , No 0.5mg/l N, c 0/No 1,62 ,,, .5-lt.5

Drea.n.J.•uint .Kineties , No o.s

. Pa:lln ~ s Data N, Co/No 1,62 ... 5-46

Figure;

?J ,_

27

JO

Jl

Computc~d

at pH 7~5, ed Brealcpo

pH 7, '?, 15oc, Compu

at pH

cs vs~ 0 ~ 5 rng/l N w

vs ~ Pc:tlin ~ s N, Co/No 1,62

s vs. Pal ~s ta N, Co/No l ,[-)2

vs ~ Pal ~ s 0,5 mg/1 N, Co/No 1,

Kineti.cs vs. Pal •s at pH

Compu at pH , No 0.5 mg/1 N, c0/No 1,82

c

ted Breakpoint 8 vs, Paiin' 8 pH 7,0, 15oc, No 0,5 mg/l N, Co/No 2,J7

j.cs vs. !s Data

. ' '

Computed at pH 0,5 rng/1 N, Co/No 2.3? ,,,

at K

, No

vs~ PaJ ~s Data 0. 5 mg/1. N t Go/No 2 ~ 37 @ ~

ics vs~ Palin's Data 0,_5 N, Co/No 2._3'? , , ,

X

50

5=53

Table

:c- 3 2

Lt-1.

5-2

.3

5-6 7 8

5-9 6.

i

SYNOPSIS

a y c ent and

ly purpose

of inf(H~ dely emplo 5. :n

mo:r:·'e than e sm of

its reactions far from bein.g es hed,

As a t of 1 a led and

e ed

that would ac with knowt1 of reac 8

srn had e react as

t.flesl.scy

e rate for t1;es e re-

a.(:tlo.n s use o:f numr~r al

tations with the

ass tane~~ of a

patterns for the as a whole and

i.dua1 em any

1 ·'·

The ts of such or1s .re•-r ea1 ed.

sev features of the constructed model of

reactions, :LncJ occu.rronc~e of G

max NHC12 co:nc; of ee

on model,

thE: Elffects of reac cone ent.:rat so

with rr;;s t to these

The c also showed that e

OYl

of an es

}; c

to bo a cruci

and

Moreover, a

ono for

the

chlo:cir1e to a.mmo:rd.a ra o near 1~8 Qeemed most 1 to

yi signi

Bas

comput;at

·wJ.thi:n

tune

as feasible,

data for the eluc of

e fin_dings of

studies were an_d cone

range of itions, but

t on were made as

tf1e course of the reactlon.G as many

s ies o:f con.cern as feas e vvere mea.,su.red as

y as s a tur13

of the reac

res of such studies were t.i

ace with those of ons ~

featu:r·es and ctions bar:1ed on the ts of computations

werH ~ howevt~r * were

r

lf

be us as

reac ons

es that were not reso

rates o:f c

accord with the

f ation of the

spec if

which are mos y e at the ent,

There were thret:-: major w E;

F

e

to

appears

n1easttremer1ts t

find-

the concen-

i .:i

tration of substance es e

of es a maximum after· a few

on of ne

cuss of the break"·

po reac had 011 stoi t

NHC12 was some sort of an the t

demonstrat that ern of its

growth decay that of a true tance~

S r•eondl very unusual c one of thr:

ti reactants reac , HOC1,

ti Y r at a of reac on~ a

~ was con=

~fe;;:;,t·ure of the ess

for a c

t(~ncc t.):f'

ra:rlge of c OY'l.S., Not the conf

of

o serves to er-3 ma.;jor

roac 1 to

2NHC12 + + JCl + HOCl

It t to conct~ tHrnatc:: rt=::ac t

of sto h scrt fy the

at the same t c1 retu.rn o:f free ch1or

y, the

xv

fJ.t constant on f 1

kc e at pH 7.0 to 12,85 e at pH 6,7, So

to be of tlo:n

of

'Nater E:tS V?(:01l as on.

s such as should o be

the t of

reaction wh effect of s cd

at

CHAPTER I

CHLORINATION OF WATER AND EFFECTS OF AM!I'lONIA NITROGEN

Since its introduction into water treatment nearly

seventy years ago, chlorination has become almost the only

method used for effective disinfection of water supplies. In

the foreseeable future, this situation appears likely to per-

sist 1 in addition, use of chlorine will increase with its

application in the field of wastewater treatment.

Because of its high oxidation potential, aqueous chlo-

rine reacts with many oxidizable chemical impurities in nat-

ural waters and wastewaters. This results in a loss of ac-

tive chlorine which is commonly called the "chlorine demand"

of the water. "Chlorine demand" represents chlorine-wasting

side reactions that must be satisfied over and above the con-

centration.of active chlorine required for disinfection.

The presence of ammonium in natural waters is ubiquitous,

and its reaction with aqueous chlorine plays a role much more

important than "chlorine demand" per se. Two of its distinc-

tive features are the formation of chloramines--commonly

called "combined available chlorine"--and the interesting

phenomenon of "breakpoint" that develops under proper condi-

tions of chlorination.

A. Chlorination of Water

"Chlorination" has been almost synonymous with "disin-

fection" in the field of water treatment, especially in the

1-1

u.s. This predominant position has been gained because of the effectiveness of aqueous chlorine as a germicide, its

ease of application, measurement and control, its freedom

from toxic effects, and its economy for large-scale water

works,

When elemental chlorine is dissolved in water, the

following reactions occur:

Hydrolysis=

Cl2 + HzO = HOCl + H+ + Cl-Ionizationl

HOCl = H+ + 001-

(1-1)

(1-2)

The hydrolysis constant is of such magnitude that no

measurable concentration of 012 remains in solution when the

pH of the chlorinated water is more than about 3.0 and the

total chloride concentration is less than about 1000 mg per

1. At ordinary water temperatures the hydrolysis of chlorine

is essentially complete within a few seconds, and the ioniza-

tion of HOCl produced is in essence an instantaneous, rever-

sible reaction.

Like most treatment processes in sanitary engineering,

chlorination of water became a standard practice long before

scientific knowledge of the chemistry was well established,

Chlorination of public water supplies was started roughly in

the first decade of the twentieth century, but the acid ioni-

zation constant of HOCl was not well established until 1935·

Once it was recognized that HOCl is much more effective

1-2

than OCl- as a disinfectant, the importance of their dis-

tribution in aqueous systems became apparent. Operationally,

however, these two species were lumped together as "free

available chlorine". As far as disinfection capacity is con-

cerned, this classification can be confusing and misleading,

because a measurement on "free available chlorine" alone does

not tell exactly how much germicidal power there is in the

system. Nevertheless, the maintenance of free available chlo-

rine in treated water supplies certainly serves as an assur-

ance of the hygienic safety of the water. The practice is

commonly called "free residual chlorination,"

B. Ammonia in Natural Waters

Ammonia in natural waters is an important chemical

species in the nitrogen cycle (1), Primarily, it results from

the decomposition of various N-containing organic compounds,

such as proteins from plants and animals. For example, mam-

malian urine contains the nitrogen resulting from the metabol-

ic breakdown of proteins. The nitrogen exists in urine prin-

cipally as urea, which can be hydrolyzed rapidly by the enzyme

urease to ammonium in an aqueous environment.

/Hz \"' o + 2H2o (1-3)

NH2

Generally, the ammonia in polluted waters has been in-

troduced from agricultural runoff, from sewage effluents, and,

because of the high solubility of ammonia in water, from rain

1-J

water. In addition, as noted above, ammonia can be released

by decomposition of organic nitrogenous materials in the

water.

The nitrogen in ammonia has an oxidation state of -J,

the most reduced form. A most important chemical property

of aqueous ammonia is its weak basicityt + NHJ + H20 = NH4 +

( NH4

+) ( OH- )

(NH))

OH-

(1-4)

At neutral pH, most of the ammonia is in the cationic form,

while at a pH of about 9.6, the distribution is about equal

between the un-ionized and the cationic forms(l).

At the pH values characteristic of most natural waters,

ammonia is converted to other nitrogenous compounds mainly

through biological mediation. Under aerobic conditions, am-

monia may be oxidized to nitrite and nitrate by autotrophic

nitrifying bacteria, which obtain energy from this oxidation

reaction, Nitrate, in turn, may be used by plants for con-

version to organic nitrogen. Alternatively it may be reduced

to nitrite and then to nitrogen gas under anaerobic condi-

tions, in the presence of an organic energy source, by hetero-

trophic bacteria. All of these transformations can affect the

concentration of ammonia in natural waters.

C1 Formation of Chloramines

The presence of ammonia in a water has a profound ef'.fect

on the results of chlorination. The formation of chloramines

1-4

is a direct result, a process that has been found to proceed

in three stages(2)1

NHJ+ HOCl = NH2Cl + H20 monochloramine (1-5)

= NHC12 + H20 dichloramine (1-6)

NHC12 + HOCl = NCL~ + H20 nitrogen trichloride (1-7)

or trichloramine

The conditions that limit the formation of the three

chloramines in aqueous solution were extensively studied by

Chapin (3, 4), The products of the reaction between aqueous

chlorine and ammonia show a high degree of dependence on ex-

perimental conditions. When the reaction was effected through

rapid mixing of sufficiently cool (at room temperature) and

dilute buffered solutions, with ammonium ions always in ex-

cess, the nature of the products was found to be governed by

the pH cof the reacting system()), Chapin's results indica-

ted that, under these conditions, at pH greater than 8,5,

only monochloramine is formed; in the pH range 4.4 to 5.0,

dichloramine is formed; and nitrogen trichloride is formed

only at pH less than 4.4. Between pH 5.0 and 8.5 mono- and

di-chloramine coexist in ratios fixed by the pH of the system,

At pH 7 approximately equal amounts (as chlorine) of mono-

and di-chloramine occur.

Nitrogen trichloride was found to have the strongest

odor, volatility from aqueous solution, and relative solubil-

ity in various immiscible solvents, such as carbon

1-5

tetrachloride, chloroform, and ether. In these properties,

dichloramine was reported to be intermediate between nitrogen

trichloride and monochloramine.

Unfortunately Chapin's results, which have been widely

quoted in the literature of sanitary engineering, are not

valid for normal conditions of water chlorination. They

appear to represent the equilibrium condition, which is

readily established for relatively concentrated solutions.

Chapin used concentrations of total aqueous chlorine in the

order of 2,000 milligrams per liter, while the concentration

normally used in water chlorination is only a few milligrams

per liter. The several hundredfold difference in concentra-

tion affects greatly the rates of reactions involved in

equilibration,

Conceptually, the effects of concentration on the chem-

istry of chloramines was recognized by Berliner in his review

article dated 1931(5), only two years later than Chapin's

work()). After a comprehensive review of the chemistry of

the chloramines, Berliner concluded that they exhibited un-

usual physical and chemical properties which were highly de-

pendent on their aqueous concentrations, and that a complete

separate chemistry could be written for them in various

ranges of concentration. He also indicated that concentra-

tions in the order of a few milligrams per liter were of most

importance to those interested in water treatment. Further,

Berliner stated, "We do not have for these conditions as

:1.-6

complete a knowledge as desirable of the rates of reaction

and conversion, the effects of mass action of the reactants,

or the influence of sulphates, carbonates and other such

materials as are found or are added to water under treatment,"

In spite of Berliner's insight of the early 1930's, the

importance of concentration and the role of chemical kinetics

was not fully appreciated and studied until the early 1950's,

Instead, discussion was focussed upon the distribution bet-

ween mono- and di-chloramine in chlorinated waters when the

molar ratio of chlorine to ammonia is one or less. Knowledge

of this distribution is essential for evaluation of overall

disinfecting efficiency in the chlorinated water containing

excess ammonia. Just as the disinfecting efficiency of

solutions of free available chlorine will vary with pH

because of the change of distribution between HOCl and OCl-,

that of chloramine solutions may be expected to vary depend-

ing on the distribution between NH2Cl and NHC12 •

In interpreting the results from Chapin's work, Fair and

coworkers(2) proposed that equilibrium existed in accord with

the reaction:

(1-8)

They accordingly performed additional experiments with

varied ratios of chlorine to ammonia, measuring the chlora-

mines spectrophotometrically and an equilibrium constant was

computed from the results. Based on calculations with this

equilibrium constant, the distribution between NH2Cl and NHC12

1-?

was computed for various conditions. Probably because of the

low molar absorptivitics of the chloramines, the concentra-

tions used experimentally were still so large that some

equilibration presumably took place subsequent to the forma-

tion of the chloramines. The results therefore have no gen-

eral validity.

Palin's extensive study (6, 7) was a successful attempt

to provide systematic data on the formation of chloramines

under the general condltions of water chlorination. An

important feature of his study was the use of differentiating

procedures. Stepwise colorimetric titrations, utilizing

neutral o;rtho-tolidine as indicator and ferrous ammonium

sulphate as titrant, were devised to differentiate not only

free chlorine from combined chlorine, but also monochlora-

mine, dichloramine and nitrogen trichloride, i.e. various

components of combined chlorine. For easier reference later,

this analytical procedure may be called the NOT-FAS method.

A wide range of molar ratios of initial chlorine to

arrw.onia, from 0.21 to J,16, was covered in Palin's work(?).

He found that with equimolar initial chlorine and ammonia

about 50% of the total chlorine was converted to NHC12 at

pH 5 and only about 5% at pH 7, much less NHC12 than pre-

dicted from Chapin and from Fair ~ ~. He also demonstrated

variation in the fraction of chlorine converted to NHC12 with

changes in the molar ratio of initial chlorine to ammonia.

The initial concentration of ammonia-nitrogen was kept at 0,5

1-8

milligram per liter for all his experiments, and the temper-

ature of his experimental system was about 15°C.* Temperature

effects were not studied.

Although Palin's work was rather extensive, covering a

broad range of pH values and of ratios of initial chlorine to

ammonia, his data are very incomplete kinetically, for he

gave results only for a few times of reaction--ten minutes,

two hours, and one day--and for a single concentration of

ammonia. Conventional kinetic analyses to find the order of

reaction and specific reaction rates cannot, therefore, be

performed on his data. The only kinetic data, until now,

came from the studies by Weiland Morris (8,9).

The kinetics of the formation of monochloramine was

thoroughly studied by Weil and Morris(8). The reaction rate

was found to be first order with respect to each reactant

according to reaction (1-5). The reaction proceeded quite

rapidly and went substantially to completion in about one

minute under ordinary conditions of water chlorination. How-

ever, the rate was very much dependent on the pH of the solu-

tion, The maximum rate occurred at a pH of 8.4, and the rate

decreased rapidly at greater and lesser pH values. It was

proposed that the actual reactants are the neutral molecules,

HOCl and NH3• The variation in rate with pH could. then be

calculated precisely on the basis of the solution equilibria

*Palin, A.T., Private communication,

1-9

of the reactants. The Arrhenius activation energy was found

to be 2.5 Kcal.

Wail and Morris(9) also studied the kinetics of the for-

mation of dichloramine in an indirect and interesting way.

Buffered solutions of chlorine and ammonium salt, the latter

always in excess, were mixed in various ratios, at pH ranging

from 4.5 to 6, and were allowed to react completely. Then

ultraviolet absorbance values of the solutions were measured

at 245, 260, and 29 5 mp. and the relative amounts of NH2C1

and NHC12 formed were calculated from the measurements.

The choices of acid pH values and an excess of ammonia

were required both to prevent loss of chlorine through the

decomposition of dichloramine, and to give conveniently

measurable quantities of both mono- and di-chloramine. Fur-

thermore, the presence of excess ammonia insured that no un-

reacted HOCl would remain in the solution.

The formation of dichloramine, like that of mono-chlora-

mine, was found to be first order with respect to each of the

reactants as shown in reaction (1-6). The overall reaction

process, i.e. reactions (1-5) and (1-6), thus constituted a

competitive, consecutive reaction scheme. The equations for

the distribution of products in such a reaction system were

thus applicable. After reaction of all of the free chlorine

in the experimental system(9), the final solution would con-

tain NH2C1 and NHC12 in a ratio determined by the specific

rate for reaction (1-5), k1 , relative to that for reaction

1-10

(1-6), k2 , and by the initial ratio of free chlorine to

ammonia. Therefore, the ratio, k1 1 k2 , could be calculated

from the distribution of reaction products, NH2Cl and NHC12 ,

determined spectrophotometrically. k2 could thus be evaluated

by use of k1 (9,10), which had been determined independent-

ly{ 8) •

One of the complicating features of the study was the

finding that the rate of formation of dichloramine was sub-

ject to general acid catalysis, being dependent both on

hydrogen-ion activity and on concentration of acetic acid,

the buffer used in the study by Wail and Morris(9). These

effects were eliminated by standard techniques.

Based upon the specific rates for the formation of mono-

and di-chloramine, Weil and Morris(9) calculated the distri-

bution of NH2Cl and NHC12 as a function of pH, temperature,

and molar ratio of chlorine to ammonia. In agreement with

Palin(?), the fraction of the chlorine converted to NHC12 at

pH 7 and with equimolar initial reactants was indeed quite

small, amounting to but 6% at 20°0 and 13% at 0°C, as com-

pared with Chapin's value of 50%(3), This difference was

attributed to different natures of the reaction processes at

the different levels of concentration, Chapin's experimental

conditions, i.e. high concentrations of reactants and buffers,

favored equilibration of the chloramines. His result might

therefore represent the equilibrium condition of the system.

On the other hand, the distribution of chloramines obtained at

1-11

the low concentrations used by Palin and by Weil and Morris

was not a mixture in equilibrium, but the composition deter-

mined by rates of formation of the individual chloramines,

It appears then that dichloramine is of little importance

in the chlorination of ammonia-containing waters with neutral

or alkaline pH unless an excess of aqueous chlorine over an

equimolar ratio of initial chlorine to ammonia is used.

When a molar excess of initial aqueous chlorine over

ammonia is used, the formation of dichloramine may become

significant, and the interesting phenomenon of "breakpoint"

can be observed,

D. Breakpoint Chlorination

About 40 years ago, an interesting phenomenon was first

observed during the treatment of water by chlorination. As

the chlorine dose to some waters was increased, the chlorine

residuum, obtained after a period of contact, behaved in an

erratic manner, i.e. a high chlorine dose gave a lower resi-

duum than was obtained with a low dose. When chlorine residua

were plotted systematically against the chlorine dose, plots

like the curve shown in Fig. 1-1 were typically obtained.

A

Residual

Chlorine

Dose of Chlorine

Fig. 1-1 Traditional Breakpoint Curve

1-12

Initially as the chlorine dose was increased, the

residual chlorine, measured normally after a contact period

of one to several hours, also increased until a "hump" was

reached, shown as point A in Fig. 1-1. Beyond this point,

the residual chlorine decreased with increased dose of

chlorine until point B--called the breakpoint--was reached.

Any further increase in chlorine dose resulted in a propor-

tional increase in residual chlorine, mostly in the form of

free chlorine, It was also established that the chlorine

dose required to reach the breakpoint was closely related to

the concentration of ammonia-nitrogen in the water to be

chlorinated.

A more quantitative and significant diagram of break-

point chlorination is shown in Fig. 1-2 AI the significance

of some of the details will be d.escribed later.

The discovery of the breakpoint phenomenon led to a

uniquely convenient and effective form of controlled chlor-

ination called breakpoint chlorination. Based upon the lo-

cation of breakpoint, which was determined by laboratory

experiments, proper chlorine doses could be used just to

exceed the brealcpoint and leave any desireable level of free

residual chlorine. By doing so, great improvement could be

effected on the efficiency of disinfection and the palata-

bility of the treated water.

Because of its practical importance, breakpoint chlor-

ination has been studied by numerous investigators. Only

1-13

Fig

ure

1-2

S

CH

EM

AT

IC

BR

EA

KP

OIN

T

CH

LO

RIN

AT

ION

D

IAG

RA

M

( B

ased

up

on

Pali

n's

D

ata

at

pH

7,3

an

d aft

er

1 D

ay

of

Reacti

on

)

Ci

No

1

mo

ls

resid

ua

l o

xid

izin

g

ch

lori

ne

p

er

mo

l a

mm

on

ia

A

NH

2C

I tr

ace o

f N

HC

l2

Fre

e

Ch

lori

ne

0!'

--~

I

No

mo

ls

resid

ua

! n

itro

ge

n

per

m

ol

am

mo

nia

1 2

3

1

c

H~- "'

Ol-

-~

1 ~

~

0 1

2

("'!

"-

"'lO

No

mo

is

oxid

izin

g

ch

lori

ne

a

dd

ed

p

er

mo

l a

mm

on

ia

3

Fig

. 1

-2

A

Fig

, 1

-2

B

.... I ~

those most significant and most relevant to the present re-

search will be reviewed in the following paragraphs.

Unaware of any practical applications in water works,

Chapin(4) observed phenomena that agreed closely with what

was later to be called the breakpoint in his pioneer studies

on the formation and decomposition of the chloro derivatives

of ammonia. When buffered solutions of chlorine and ammonia

were mixed in a molar ratio of initial chlorine to ammonia

in excess of' unity, loss of residual chlorine was observed

at pH of 9. 0 and 4-, 9, at 25°C. The molar ratio of chlorine

lost to ammonia lost was found to be about 1. 7 for both pH's.

The initial concentration of ammonia-nitrogen used by Chapin

was about 70 milligrams per liter, however, much higher than

those encountered in water treatment.

Qualitatively, Chapin observed that at pH 9 the mixtures

evidently ran through rapid alternate formations of chloro

derivatives and decompositions of the latter under the in-

f'luence of' hydroxyl ion until either ammonium ion or hypo-

chlorite was exhausted and only monochloramine and ammonium

ion, or hypochlorite, remained, action being found already

practically complete in one-quarter hour.

He also stated that at pH 5 the mixtures ran through

a similar series of decompositions up to a molar ratio of

close to 1.7, but here excess of chlorine evidently inhibited

the decomposition of nitrogen trichloride, and the residual

chlorine beyond the breakpoint was found to be higher at pH 5

1-15

than at pH 9.

Chapin also observed that when the reaction was effected

with much higher concentrations at pH 5, the molar ratio at

the breakpoint was about 1.5 instead of 1.7. He proposed that

the major stoichiometric reaction was

(1-9)

This reaction accounts for the observed stoichiometry at the

highest concentrations. Departure from reaction (1-9) ap-

peared to increase with dilution, and could be roughly ac-

counted for by assuming the formation of nitrate and nitrous

oxide,

A systematic study on breakpoint chlorination was

probably first done by Griffin and Chamberlin(11). It was

Griffin who devised the term "breakpoint." A series of' ex-

periments was undertaken to study the effect of chlorine on

water containing 0.5 milligram per liter of ammonia-nitrogen

at pH 5, 6, 7, 8, and 9. Throughout the work ammonia-free

distilled water with adequate buffers was used. The tem-

perature of the system was held between 7 and 9°C. Total

chlorine remaining at the end of 20-min., 2-hr., and 24-hr.

intervals was determined by titration with iodidestarch-

thiosulfate. Total nitrogen at the end of 2-hr., and 24-hr.

intervals was determined by distillation and Nesslerization.

The molar ratio of chlorine to ammonia at the breakpoint

was found to be about 2,0, Qualitatively, there were indica-

tions that the residual chlorine within the hump part of the

1-16

curve was chloramine (mostly NH2Cl at pH 9.0 and NHC12 at

pH 5.0), and that beyond the breakpoint it was mostly free

chlorine. Generally, there was no measurable loss in ni-

trogen until the chlorine dose was high enough to reach the

"hump." From this point on, the nitrogen decreased relative-

ly as sharply as the chlorine did when chlorine dose was in-

creased, and both were exhausted at the breakpoint.

Based upon limited kinetic observations at their three

intervals of time, the overall rate of the breakpoint re-

action was found to be highly dependent on the pH of the

solution. The breakpoint was achieved most rapidly at a pH

between 7 and 8.

Marks and Glass(12) made some interesting observations

on breakpoint chlorination. In the course of their develop-

ment of amperometric titration as a new analytical tool for

the determination of residual chlorine. Qualitatively, they

confirmed the existence of nitrogen trichloride in the course

of breakpoint chlorination, However, the concentration of

nitrogen trichloride seemed to be rather small until the

chlorine dose was equal to or greater than the molar ratio

to N for the breakpoint.

Moore, Megregian, and Ruchhoft(13) studied breakpoint

chlorination using an approach similar to that of Griffin and

Chamberlin(11), except that more discriminatory analytical

methods were employed. The initial concentration of ammonia-

nitrogen was maintained at 0. 5 milligram per liter by adding

1-17

ammonium chloride to chlorine-demand-free water, buffered at

the desired pH of 6, 7, 8, or 9. After a proper amount of

aqueous chlorine was added to give the desired dosage of

chlorine, the mixture was allowed to stand for contact times

ranging from 2 to 24 hours1 then a number of analytical de-

terminations were applied to the samples, These included!

total residual chlorine by the ortho-tolidine method, free

residual chlorine qualitatively (flash test) and quantita-

tively by n-aminodimethylaniline method, free and combined

residual chlorine by amperometric titration, oxidation-

reduction potential, and total nitrogen by Nesslerization.

The temperature of the system was about 25°C.

The molar ratio of chlorine to ammonia was found to be

about 1.0 at the hump, and about 1.8 at the breakpoint.

Essentially the same type of breakpoint curve was observed

throughout the pH range studied.

For each breakpoint curve, no chloramine was found be-

yond the breakpoint, and no free chlorine was found until

breakpoint was reached except at low pH. At pH 6.0, the flash

test showed dubious existence of free chlorine under the hump,

especially between the hump and the breakpoint, and this phen""

omenon persisted throughout the contact period from 2 to 24

hours, Consequently, it was speculated that a partial

reversal of the chlorine-ammonia reaction might occur at the

lower pH, which would result in the liberation of free chlo-

rine. The decomposition of monochloramine was suspected

1-18

to be the key reaction.

As summarized in the previous section, Palin's extensive

study (6, 7) also covered the decomposition of chloramines

and, consequently, breakpoint chlorination. In addition to

the NOT-FAS method described before for discrimination of

forms of active chlorine, Palin also used direct Nessleriza-

tion to determine free ammonia-nitrogen, and direct Nessler-

ization after dechlorination to determine the total of

ammonia-nitrogen and chloramine-nitrogen. He was able with

these analytical tools to illuminate the breakpoint phenom-

enon in greater detail. Based upon his data with modifica-

tion, Figure 1-2 is presented as a schematic diagram of

breakpoint chlorination at neutral pH, showing the detailed

chemical composition after complete breakpoint development.

Palin concluded that the reactions between chlorine (as

HOCl) and ammonia led to the formation of monochloramine,

dichloramine, or nitrogen trichloride, or to some mixture of

these compounds. Formation of the more highly substituted

derivatives was favored by increased acidity of solution and

increased molar ratio of initial chlorine to ammonia. Some

combinations of products were found to be stable at certain

limited pH ranges, e.g. NC13-c12 , NHC12-NH3

, NH2Cl-NH3, and

NH2c1-NHC12-NH3. Generally, decomposition occurred if there

remained an excess of chlorine after chloramines were fanned,

At the breakpoint, an essentially complete oxidation-

reduction process occurred, leading to the disappearance of

1-19

all the ammonia and chlorine from the solution,

The molar ratio at the breakpoint was found to be 1,88

at pH 6, 1.62 at pH 7, and 1,68 at pH 9. The final products

of the overall reaction were proposed to be predominantly

nitrogen plus relatively small fractions of nitrate. The

proportion of nitrate increased with increased molar ratio

of initial chlorine to ammonia. In two experiments at neu-

tral pH and with molar ratios of 1.82 and 2.25, Palin was

able to account quantitatively for all of the chlorine added

initially by assuming nitrate as the sole by-product in

addition to the major product of nitrogen. The formation of

ni.trate was postulated to be connected with the decomposition

of dichloramine.

Nitrogen trichloride was formed most readily at pH 4 or

less, but might be found at pH values up to 8 when the molar

ratio of initial chlorine to ammonia was sufficiently high.

Beyond the breakpoint, the nitrogen trichloride formed was

quite stable in the presence of excess free chlorine.

Based upon these and other experimental observations,

some detailed chemical reactions were also suggested to

account for the stoichiometry. These will be presented and

discussed in Chapter II.

Morris and Weil (14) conducted kinetic studies on the

decomposition of dichloramine on the assumption that this

was a key reaction in the breakpoint phenomenon, The rate

of this decomposition was found to be first order with respect

1-20

to NHC12 and inversely proportional to the hydrogen•ion

concentration in the pH range of 6.5 to 9.0. It was found,

moreover, that the rate of decomposi t.ion and the amount of

chlorine lost were not affected by the presence of NH2Cl.

One disturbing complication was the finding that the decom-

position of NHC12 in acid solution was greatly accelerated

in the presence of HOCl. Attempts to study t.he accelerated

process kinetically were not successful.

A comparison between the rates of formation and the de-

composition of NHC12 shed new light on a possible mechanism

for breakpoint development. At pH values less than about 7.5

and with concentrations of a few milligrams per liter, the

decomposition of NHC12 was slower than its rate of formation!

at pH values greater than 7.5 the rate of decomposition was

faster than the rate of formation. Morris and Weil con-

cluded, therefore, that the rate of the overall oxidation-

reduction reaction of the breakpoint process was controlled

by the rate of decomposition of NHC12 at pH less than about

7.5 and by the rate of formation of NHC12 at pH greater than

7.5.

Kinetic experiments on the overall breakpoint reactions

were conducted in the pH range of 7.7 to 10.0, with the molar

ratio ranging from 2,0 to 4.0. Assuming the formation of

NHC12 as the rate-limiting step in the alkaline range, the

rate of breakpoint development could be evaluated from the

measured rate constants for acid solutions. The observed

1-21

rates of loss of active chlorine were found to be in ex-

cellent agreement with those calculated from the rate of

formation of NHC12 (9), Similar kinetic experiments in the

acid range yielded qualitative agreement, but no quantitative

treatment was possible because of the unknown magnitude and

pa:t;tern of the acceleration by HOCl on the decomposition of

NHC12•

A detailed reaction scheme that will be presented and

discussed in Chapter II was proposed for the mechanism of

the decomposition of dichloramine.

E. Purposes of This Study

With the termination of the work of Weil and Morris

in the early 1950's, detailed investigations of the course of

the breakpoint reaction between chlorine and ammonia virtually

ceased. As a result even though breakpoint chlorination has

been widely employed in water and wastewater treatment for

more than thirty years, the mechanism of its reactions is far

from being established. This deficiency represents not only

a lack of the theoretical understanding, but also a great

loss in the ability to manipulate water chlorination con-

structively for maximum efficiency and benefit.

The present research has been undertaken to develop a

mechanistic model for the breakpoint process, which can be

simulated by computer, tested by laboratory experiments, and

used to predict the pattern and results o.f breakpoint

1-22

chlorination over a wide range of conditions, As the demand

for more efficient disinfection grows, it is hoped that this

mechanistic model can serve as a basic scientific tool for

the optimization of breakpoint chlorination process.

1-23

CHAPTER II

DEVEWPlVIENT OP A MECHANISTIC MODEL

A. General Characteristics of Complc:x Reaction Scheme

From facts that the stoichiom•o>try of the breakpo

2

ess not simple, that the stoichiometry unchanged

despite the e format of as a function

of pH, and that the development of the breakpoint a

maximum in tile pH range 7 to 7.5, it is evident

the breakpo chlorination

of several elementary react

a complex reaction compos

Generally, a complex

reaction scheme may consist of various combinations of

el

back

reactions,

ons,

as consecutive,

ion of the pattern

:for complex react

Complete det

Hke these

ion of the

simultaneous

solut kinetic

each of the individual steps.

As illustrated Frost and Pearson(16), intermediate

species

during

cons reactions ordinarily go through maxima

course of reaction. The positions and magnitudes

of the maxima will depend on the ion orders and the

values of the specific rates for the formation and

consumption of the species.

important charac c of consecutive reactions

the outstanding role of the "rate-limiting" step or steps

the overall reaction pattern. The overall reaction rate

is ess ly determined by that of the rate-1 step

or s establ equilibria or reactions leading up

to the slow steps may also be important for setting the

reaction order and absolute rate. Any reactions subsequent

to the rate-limiting step will not affect the overall

reaction pattern or kinetics, however, although they may

affect the stoichiometry of the overall reaction. As a

consequence, kinetic studies, like stoichiometric relations,

are powerless to give information about any reaction

mechanism following the slow steps,

It must also be noted that the "rate-limiting" step or

steps are not necessarily the same under differing reaction

conditions, Specific rates of reactions are affected in

various ways and extents by changes in temperature, pH or

other reaction condi ons. Hence an elementary reaction that

rat imiti.ng for a particular mechanism under one set of

conditions may be replaced by a totally different one in

other circumstances,

B .. PreviouslLJl;:_ollos ed Mechanisms

Previous investigators, except for Well and Morris, have

based their considerations of mechanism for the breakpoint

process largely on stoichiometric information and occurrence

of chloramine intermediates, for this was all that was avail-

able to them. Their formulations of mechanism have, in

consequence, been sketchy and incomplete. In spite of this

situation and spite of the fact that any mechanistic

reasoning which does not take kinetics into account must be

quite hypothetical, there are significant clues in the

discu.ssions. So they wi be reviewed some detail.

2-2

Four significant attempts at mechanistic interpretation

may be noted:

L R. M, Chapi f1• As presented in Chapter I, Chapin( 4)

proposed the following equation as the stoichiometric

reaction for breakpo chlorination:

( 1-9)

Chapin( Lf) also proposed two general hypotheses to account

for his numerous observations on the format and decom-

position of the chloro derivat of ammonia. 'I' he

was that there is a hydrogen-ion-induced hydrolysis of each

chloramine to NH4 HOCl, particularly below a characte-

pH, that resulting HOC1 reacts with

unhydrolyzed chloramine to produce a more highly chlorinated

derivative. 'I' he second hypothes was that hydroxyl

induces formation of Cl- icularly

above a charac pH, with gaseous nitrogen and OC

as the principal associated products. The first hypothesis

accounts for the formation of NHC12 and NC1 3 through the

ac fication of NHzCl and NHC1 2 , ; but no

of able involved in e and,

, the hypothes to do with

ion. It may noted o that neither

,, NHC + H' + 2H2o

+ NH4 + 2HOC1 ( 2-1)

nor NClJ + H+ + 3H2o .k

NH[f + 3HOC1 (2-2)

can possibly be regarded as an ementary reaction,

2-3

The second hypothesis was deduced the

decomposition NHClz and NCl3 with increase of pH.

Thus it might be significant for the breakpoint reaction,

Chapin concluded that the principal reactions for

decompos of NHCl2 and NCl3 were:

2NHC1 2 + 4-0JC --- JCl + OCl + }H2o + N2

2NClJ + 60H- JCl i· JOCl + 3Hz0 + N2

(2-3)

(2-1~)

'rhe residual it ion

presence of excess

t at pH greater than 9 and

was identified as being

primarily NHzCl, apparently as a result of the react

en the l OCl and excess ammonia, Neverthe-

less, Chapin reported experimental evidence the evanes-

cent existence of 1

Again, however, except for the

OH-, neither of these

mechanism,

At molar l .67 of

has anything to do with

tia1 ehlorine to ammonia

(0.005lVI) and at pH 5, a gaseous produet was reckoned as

of

nitrous oxide, and it could quantitatively aeeount the

departure from reaetion (1-9). But at pH 9 no nitrous ox

was found at mo ratios 1, 60 or 1, 67,

s e the

various decomposi

suggested that the ni

of nitrous oxide resembles the

of hyponitrous ac , H2N202 , Chapin

group, HNO, could be an

mediate in ( 2-4) •

2, ,J, R. JSossum. Rossum(l5) proposed as a meehanism

breakpoint chlorination that NHzCl NHCl2, after

2-4

being formed, react as follows:

NHzCl + NHClz Nz + 3HC1 (2-5)

With the aid of mass action equations for the formation of

NHzCl and NHC12 , Eossum calculated the variation of

residual chlorine with chlorine dosage, and obtained a result

having a close n;semblance to the breakpoint curve,

Although Hossum's idea appears plausible, it has no real

validity. F t of , the proposal was based on mis-

conception that Chapin's data on the distribution of the

was valid for low concentrations with excess

chlorine. Apart from this the written reaction again

wholly stoich c except for the suggest that NH2Cl

and NHC1z are jo

given again cannot

involved as reactants. 'rhe reaction

ibly be an ementary reac

J, A. T. Palin. Palin(?) suggested the following

reactions to account for ex tens

tions on the oxidation of ammonia by

(i) At pH 8 or greater, Palin obs

experimental observa-

chlorine.

although

excess and NHzCl might co-exist some , there

was a gradual of NHzCl, and tended to disappear

completely when the initial chlorine excess was sufficient.

Intermediate formation of NHClz and NCl3 was not obs in

these alkaline solutions~ So suggested react was

2NHzC1 + HOCl (2-6)

(ii) Systems in which the total chloramines ially

(i.e, at 10 minutes) were almost entirely NHC were unstable,

The decomposition of the NHCl2 was accompanied by an

crease in free chlorine although there was a considerable

loss of available chlorine, The reaction suggested

was:

N2 + HOCl + 3HC1 (2-7)

(iii) Systems which contained both NH2Cl and NHC were

also unstable, and there was fairly rapid decomposition with

eventual disappearance of NH2Cl or NHC1 2 , depending upon the

concentrations at start, The presence of

excess NH! was reported to have a retarding on

decomposition, The reaction involved might

NH2Cl + NHCl2 --- N2 + 3HC1 (2-8)

(iv) In the absence of o chloramines systems

containing free chlorine and NCl3 were fairly stable, The

relative proportions were determined by pH of the solu-

tion and appeared also to be affected by NHCl2 when present.

'Phese obs could be accounted for by:

NHCl2 + HOCl (2-9)

As in previous tanoes none of ents

any

breakpo

attempt to de mechanism or the pattern of the

reaction, None of the written r can

ibly be elementary steps in a mechanistic SN''""" All

that has en done to express the desired stoichiometric

production of N2 terms of all ible

4. J.C, Morris & I.W. Weil. As a t of

kinetic studies on the decomposit of NHCl2, Morris and

2-6

2-7

Weil(i~-) concluded that this the ical reaction in

breakpoint reaction. A detailed scheme as follows was proposed

as the mechanism for this decomposit

NHCl2 -- H+ + NC (very t)

NC12 (mC) N-Cl + Cl (slow)

N-Cl + OH -NOH + Cl- (fast)

2NOH H2N202

H2N202- N20 + H20

This mechanism for the on of 2 me

chlorine for every mole of ammonia oxidized, which

(2 0)

( 2-11)

(2 2)

(2 J)

(2-14)

of

greater

than the average of the ratios found by other workers - about

1, 7 to 1, It was suggested this discrepancy might

by inclusion of the following reac as a

of the mechanism:

( 2-15)

Also, because there had been observations that the

of chlorine reduced to ammonia oxidiz eases with the

excess chlorine present, Weil and Morris suggested in-

clus of the following reaction:

(2-16)

The NO formed in this would then react with oxygen

the water to give N02 and eventually nitrite and nitrate.

Much of this scheme hypotheti and not

in any detailed sense, For example, R ( 2-11)

described as the slow step and so evidences for the detailed

reactions following this cannot be obtained from kinetic

No e the occurrence of

2-8

NOH, or of hyponitrous acid, H2N2o2 , was reported, The

decomposition of hyponitrous acid to N20 and water a well-

known reaction however, Reactions (2-15) and (2-16) are

sto om 1 they could not be elementary mechanistic

react

c. Further A,nalysis of Palin's Data

In the course of devising a mechanist model for break-

int react , an attempt was made to exhaust Palin's data(?),

Some are shown Figs, 2-1 through 2-4,

Fig, 2-1 shows fractions of to available chlorine

remaining as a function of pH for three reaction periods, the

upper curves being ten minutes of reaction, the lower ones

two hours of reaction, and the zontal 1 ines for twenty-

four of reaction. The solid 1 refer to mixtures

with a chlorine to ammonia molar ratio of 1,62, the dotted

lines to those with a ratio of 1.82, As expected, the

rate of disappearance of total lable chlorine (i.e, the

rate of the breakpoint reaction) is greatest near pH 7,2 and

decreases sharply with decrease or ease of' pH.

Figs, 2-2 through 2-4 show a detailed breakdown of the

constituents of the residual chlirine at a reaction of ten

minutes, with molar ratios of Cl toN equal to 1,62, 1,82 and

2.3?, • . 1 1.ve .•. y, Perhaps the most interesting aspect of

e results the way the concentration of NHC decreases

with increasing pH and is practically zero at pH greater 7,5,

This is in good agreement with the theory proposed by Morris

and Weil(l4) that the ion of NHC becomes rat imi ting

Q) >:: •rl

:: 0

•rl

rl

H

..C:

0 0

rl ..c:

Q)

0 rl

p .-

-l (1

j (1

j rl

;l

·rl

"d

(\j

•rl

:> w

<

(!

)

0::

:>,

rl

rl

rl

(\j

(\j

+' l·

rl

0 +

' 8

·rl >:: H

t !

_A

l;

-~--Q-

I ~

(~ /

'-..

v t=

1

0 m

w.-

I

.6'

\ ~

Ct

Co

l'\ ~\

1 ,~

;~-

~,~,:·

,,,,,-

. '

\ \

"' .£.1-

, ---

\ \

' \. ' ..

......

----

----

6

I /

r I I I I I • f I

,.,~

I P

I

~;:;

~ I

&?~I

/

/ /

..fr

'-_,.

t

=

2 h

rs ·

/

--'f

Mil

f·\''>

•'A I

'~v(Jl,

',Ea.

/ ,

I o

I ,..

QS'

, 1

.62

r

---

No

I . ,

I

I I _

Co

, ---

No

1-I

:i:)

,-0

,

,/

~

" --

-·--

. "''

b ~"'

c:::s:

i) "t

-:;; '

24

-hrs

--".l

:>',~p

·, , ,,

,-''

, (

• •

I ... ~~,--~

"--)'

7 PH

8

1.8

2 9

Fig

. 2

-1

Dev

elo

pm

ent

of

the

Bre

akp

oin

t w

ith

C

hlo

rin

e to

A

mm

onia

-N

itro

gen

Mo

lar

Rat

io

Eq

ual

to

1

,62

an

d 1

.82

To

tal

resi

du

al

Ch

lori

ne are

calc

)lla

ted

fro

m P

ali

n's

d

ata

( 7

) fo

r th

ree co

nta

ct

tim

es.

N ! •O

1fJ +' ~

1J]

+' ~ (l)

::> ·f-' .,.., +' tJl ~ 0 0

GJ ~

•.-! H 0 rl ..c: Q

rl '0 40 •.-! 1J] GJ

e>::. 'H 0

,:; 0

•.-! 2.0 '.p 0 m H ~

6

---..

C/:N rafro 9.2 fo 1 bywt. Palin dafa

7 pH 8

2-11

I NH;:CI

9

Fig, 2-3 Development of the Breakpoint with Chlorine to Ammonia-Ni t:r·ogen Molar Ratio Equal to 1.82

Data are taken from :Palin (7) for ten-minute contact at about 15°C. Data points are not shown, for clarity, but are comparable to Fig. 2-2.

(!) .s 50 f.-1 0 rl ..

at pH greater than 7.5, the NHCl2 decomposing as fast as it

formed as a result of the

NHCl2 with increasing pH.

ing rate of decomposition of

Qualitatively, all of these results tend to support the

theory ed by Morris and Weil(lL>) that NHCl2 is the key

species the mechanism of breakpoint reaction.

D. Proposed Mechanism

Of the numerous studies on the breakpoint reaction, e

of Chapin(4), Palin(?), and Morris & Weil(9,14) are generally

most useful as a basis for elucidation of mechanism and are

thus used as the principal source of information the

development of a react model. Nevertheless, it may be noted

that there are no major discrepancies between the findings o:f

these studies and those of other major investigators, so that

a ion model thus devised should be

Rnr,,sent the present state of knowledge.

1. Reaction Scheme and Specific Rates

1y adequate to

'!'he proposed reaction scheme is shown in 'rable 2-1,

including eight individual reac

ific rates, k1 through ks.

with the respect

well established

that the initial steps in the overall reaction are the

success formation of NH2Cl NHC1 2 in accordance with

(M-1) and (M-2) in '!'able 2-l, Fortunately, the

specific rates of these reactions, k1 and k2, are reasonably

well known from studies of Weiland Morris(8,9).

NCl 3

Reaction (!Vl-3) is needed to account for the formation of

in quite acid solutions or in neutral solutions when

Tab

le

2-1

Bre

akp

oin

t R

ea

ctio

n

Sch

em

e-

(A)

HO

C I

+

NH

3 --

---)

> NH~CI

+

•

• N

CI 3

+

H

zO

NC

i 3

+

H20

->

N

HC

I 2

+

H O

CJ

Nl·

!CI 2

+

H

20

-N

OH

+ 2

H+

+ 2

CI-

NO

H

+.

NH

2C

I -

N2

+ H

20

+ H

+ +

C

l-

NO

H

+

NH

CI 2

--?

N2

-;. H

OC

!+·

H+ +

Cl-

NO

H

+

2H

OC

I -

NO

:!+

3H

+ +

2C

l-

k1

kz

k3

k4

ks

k6

k7

k 8

(M-1

)

( !V

I-2)

(M-3

)

(M-4

)

(M-5

)

( lV

I-6)

(M-7

)

(M-8

)

.

N ' ,_. .(::

2-15

initial ratio of hypochlorite to ammonia high, The reverse

of this reaction, Reaction (M-1+), is also needed to take

account of the generally observed fact that NCl; is stable

under ions of water chlorination excess free

present and to aceoun·t for slow decomposi t of

NCl3, Orders and specific rates of Reactions (IVI-3) and (M-It)

are not known. Reasonable estimation of them can made,

however, on the basis of Palin's data(?) show a linear

between excess free chlorine over that required for

NHCl2 and concentration of NC1 3 formed in neutral solution,

and on the is of ln acid pH range in which NClJ

beeomes a signifleant product when the molar of init

chlorine to ammonia-nitrogen less than 2.

In accord with conclusion of Well and !Viorris(14),

the omposition of NHClz is considered as a major step in

reaction scheme. WeLt and Morris(1LJ,) found that pre-

formed NHClz decomposed in aqueous solution ace to

t-order with regard to NHC concentrat

specific rate, however, was pH , being more

proportlonal to the hydroxyl-ion conc*mtration,

observations are reflectr1d

radieal, NOH, shown as a

o:f a reactive

Reac (M-5). The nitroxyl

( M- 5) , s

specles would serve

computet

well wlthout effect on

It may be noted that of the scheme

t half of React

e

y

a mechanistic way, The remainder of the scheme, which

sto

the slow decomposition of NHC12 , is s

ometrically. Further detail about this

reaction s can be obtained only from

f:icant only

of the

on about

later reaction intermediates and their chemistry, For

, .•• ~ .... , .... e, a reaction like (M-9) undoubtedly proce by way

of nitrite, N02, but until it found it would be

less to include an additional reaction step.

Three parall reactions competing for NOH are incor-

in scheme and shown as Reactions ( M-6) , ( M-7) ,

( NJ-8). These reactions are all hypoth al at

time iLic rates are not known for any of them. If

assumed.~ as seems , that of these reactions

are rapid compared with R (NJ-5), th(m their clute

rates are oft little concern. Only rates,

which will determine the stoichiometry and the

ammounts of the nitrogenous products, N2 and NO}, are

important.

point

of

Reac

alkal

(NJ-6) included to account for

solutions where there i.s no accumu1

NHClz the operation of Heac on ( M-·7),

Howflver, K6 will be assigned a smaller valuetl:lan K7,

the observed a of HOCl Palin's

indicates that (M-7) is more important when NHC is

ause

ent~

Reacti.on (M-7) proposed to be the

reaction,

or pathway for

evidence of the

of free chlor the decomposit of NHClz

2--16

2-17

s , however, to establish definitely the role of

R.eaction (M-7) as the major pathway determining the stoichio-

metry of the Reaction ( lVJ-8) to account

for Palin's observation

product in t.ion to

formed as an end

major pro

2, Stoichiometry of the Overall Breakpoint Reaction

Corresponding to the three pathways in Reactions ( M-6) ,

(M-7), and (M-8), the overall reactions of the respective

pathways are summarized in Table 2-2, Stoichiometrically,

Reaction (M-6) or (M-7) will lead to 1.5 moles of HOCl re-

duc per mole of NH3 oxidized, Le. a molar ratio of 1,5 at

breakpoint, while Reaction (M-8) requires a molar ratio

of 4, A propor combination of these pathways will result

in a molar of about 1,6 at neutral pH as observed in

Palin's experiments(?).

E._lechnigues of Computer Simulation

2-18

The complexity of breakpoint react and the paucity

of suitable kinetic makes it almost impossible to apply

conventional for the finding of reaction

and specific rates, and subsequently for the i-

dation of the reaction mechanism. The development of

computer technology makes possible to work in the opposite

, however. If one can a suitable

model and assign reasonable specific rates to the individual

reaction steps, numerical techniques can than be used to

compute the complete course of the reaction, Comparison of'

computed with the observed may then ei

validate choices made or indicate the types of' changes

be made to bring about agreement with the

In the proposed model, there are eight chemic ies

and consequently· eight differential involved,

Table 2-2, Reactions of Different

Pathways in the Proposed Mechanistic Model

1 . The Pathway through Reaction ( M-6) :

2HOCl + 2NHJ -- 2NH2Cl + 2H20

HOCl + NH2Cl -- NHC1 2 + H20

NHCl2 + H20 ·· NOH + 2H+ + 2Cl

NOH + NH2Cl --- N2 + HzO + H·t + Cl

3HOCl + 2NHJ ·~ Nz + JHzO + 3H+ +

2, The Pathway through Reaction (M-7):

2HOC1 + 2NH3---- 2NHzCl + 2H20

2HOCl + 2NH2Cl --- 2NHCl2 + 2H20

NHCl 2 + H20 - NOH + 2H+ + 2Cl

3Cl

NOH + NHClz - Nz + HOCl + H+ + Cl

JHOCl + 2NHJ ........._. N2 + JHzO + JH+ + JCl-

J, 'rhe Pathway through Reaction ( M-8) :

HOCl + NHJ ___. NHzCl + H20

HOCl + NHzCl --- NHC12 + HzO

NHC1 2 + H2o- NOH + 2H+ + 2Cl

NOH + 2HOC1 · - NOj + JH+ + 2Cl-4HOC1 + NH3 - NOj + HzO + 5H+ + 4Cl-

2 9

as shown Table 2-J, To serve the purposes stated above,

this set of orcHnary differential equations can best be solved

by the technique of numerical integration and by means of a

computer,

The algorithm used for the integration is

essentially a Pr tor method, shown as subroutine

Haming in Appendix I. Although NOH radical shown as a

chemical species in Table 2-J, not included the

actual computation because

a reactive intermediate, Therefore, subroutine Haming is used

t as the computer output, This output stored on tape

and subsequently is plotted on a semi-logarithm scale by a

S-C1}020 Computer Recorder as the final form of computational

results,

F. Preliminary Model Computations

1 2 General Conditions

2·~20

The iminary model computations were igned primarily

to provide a comparison with Palin's data, 'J:he spec if rate

parameters us in these computations according to the reac

scheme Table 2-1 are shown :for a number o:f pH values

Table 2-Lf. It most convenient for computation i:f the

concentrations and rate parameters are ed dim ens

less :form except for the time variable, This can be achieved

i:f of the concentrations is expressed as a. molar io

to ammonia, Consequently, parameters

·Tab

le

2-3

K;n

,,.;

,_

t:,.

...,

,";C

"S

for

Scl

-em

r->

'-i<iV11.1~

-"'1'-"'-~"'•

••

il

1 ••

._..

; a

= a

mm

on

ia,

c =

fre

e

ch

lori

ne

, m

= m

on

och

lora

mid

e,

d =

dic

hlo

ram

ide

, n

=n

itro

ge

n t

rich

lori

de

, f=

nit

roxyl,

b=

nit

rog

en

, e

=n

itra

te;

an

as

mo

lar

rati

os to

in

itia

l a

mm

on

ia,

1\!0

•

du1

=

k1 N

0 at

Ct

dt

du

2 =

k

2 N

c m

t ct

dt

du3

= k

3N

0 d

t C

t d

t

• au

4 =

k

4 n

t d

t

du5 =

k

5 d

t d

t

dur,

=

0

k6N

0 ft

m

t d

t

du

7 =

k7

1\! 0

ft

d

t d

t

du8

== k 8

N0

ft

Ct

dt

a2=

a

1 -

du

1

c2 =

c1

-d

u1

-d

u2

-du

3 +

du

4+

du

7-

2d

u8

m2

.= m

1 +

du

1 -

du

2-

du6

d2

:= d

1 +

d

uz-

du3

+ du

4-

du

5-d

u7

n2

:= n

1 +

du

3 -

du4

f 2

:::::

f 1

+ d

u5

-d

u6

-d

u7

-d

u8

b2 =

b1

+ d

u6

+ d

u7

e 2 =

c 1 +

du

8

·--------~-------·---·---···

N ' N ,_.

pH

I

k 1N

0

102

k2

N0

.10

3 k

3 N

0

104

k4

I

. 10

2 k5

I

k6N

o

k7N

o

10 k

8N

0 I

Tab

le

2-4

Fir

st

Pa

ram

ete

rs

for

Re

act

;on

S

che

me

(A

)

6.5

7 .. 0

7.

3 1.

5 7.

7 8.

0 8.

3 8.

5 --

---

0.21

0 0.

561

0.91

2 1

.19

1.

47

1.80

2.

00

2.0

3

.89

4

.75

7

.62

0

. 511

.3

99

.2

52

.144

.0

93

.358

.3

03

.2

48

.2

04

.1

60

.1

01

.058

.0

38

.1 1

I .1

1 I

.11

I .1

1 I

.11

. I

. 11

I .1

1 I

.11

.06

3

.20

0

.39

8

.63'

1 1.

00

2.0

0

3.98

6.

31

1.0

1.0

1.0

1.0

1.0

1.0

1.0

1.0

2.0

2.0

2.0

2

.0

2.0

2.

0 ·1

2.0

1

2.0

.50

4 I

.426

I

.35

0 I

.:?

.88

I .2

24

I

.142

I

.081

I

.05

4

N I rv

N

actually used are products of specific rate and initial

ammonia concentration if' the reaction is second-order, as

shown in 'l'able 2-Lf, 'l'he initial ammonia concentration for

these computations was always J,57xlo··.5!Vl, i.e. 0,_5 mg per

liter as nitrogen, the concentration used by Palin throughout

his experiments.

All the rate parameters taken from the literature were

evaluated for a temperature of 20 C, Others were estimated

to apply at the same temperature, 'rwo initial molar ratios

of' chlorine to ammonia, 1.60 & 1.82, were used for these

computations. These two molar ratios were believed to be

adequate to provide initial comparison with Palin near the

breakpoint.

2. Results and Findings

Some typical results of the preliminary model computa-

tion are shown in Figures 2-5 through 2-9, with Palin's data

plotted for comparison. Qualitatively, these results are in

general accord with what is known of the concentration-time

dependence of various ies during the course of the break-

point reaction, The rate of development of the breakpoint,

however, is much faster for the model computation than in

Palin's experiments. Nevertheless, the computations provide

some insight which was not previously expressed and appears

not otherwise attainable,

The major contribution of the computations is to show

the inadequacy of present data and to indicate the types of

2-2J

0 p

UJ (j)

·rl C) (j)

P< Ul

'H 0

0 . .,;

.2

.I

-:;; .05 -r:t::

,02

2-24

I pH 6.5 C0 / N0 1.60

Time Fig. 2-5 Results of Preliminary Model Computation (curves)

vs. Palin's Data (symbols) at pH 6. 5, Co/No=l. 60

Cf = free chlorine, lVl = mono chlorine, D = dichloramine, T = nitrogen trichloride, Ct = total chlorine

2-25

2 I

I ~ (D

pH 6.5 tuJ 0 H C0 / N0 1.82 N2 +' 1 -orl -z al

·d ~ 0

~ .5 ~

ct rl c ru •d No

.,. _.;.I +' Ct EJ3

EBI •.-l ~ cf H 0 -· +' 't =:51 Ul T CfaJ CI)I (D ---"""~~'

·d

J 0 Q) P-; .1 rJl 't-i 0

1 0 •rl t --?o +' Jt. ro .05 ~ 0 ~~ H

2

h (j)

bD 0 H +' •ri z. (\j

•rl ~ .3 0 s s -< rl (\j

•.ri +' 'H ~

H

0 c +' w (j)

•rl 0

>:: ())

b.O 0 ,, 4J •d 2;

ro •rl

>:: 0

"' ~ rl cd

·rl +' •rl

>:: H

0 +' Ul ())

•rl 0 ())

Pc rJl

'H 0

0 ·rl +' ro ::

~ .-{

0 ~

.1

c No

.031

.01

2-

9M ~Or

-

pH 8.5

- Co/No 1.60

10xN03

()D

Fig. 2-8 Results of Preliminary Model Computation (curves) vs. Palin's Data (symbols) at pH 8.5, Co/No=L6

Cf = free chlorine, !VI = mono chloramine, D = dichloramine, T = nitrogen trichloride, Ct = total chlorine

2

,::: (J)

QD 0 H +' ,,.; ?-:

"' •rl ,::: 0 8

~ rl

"' .,.; +' •rl .I ,c: H

0 _Q_ +' No rJ) (J)

.,.; 0 Q)

,03 p, (fl

'H 0

0 .,.; +'

"' .01 (}:; '. @ rl 0

2-S

.003

2-28

Nz EDCt : -l

M Ct @Cf !

-£L ---- ..., -E9 M

pH 8.5

Co/No 1.82

()D

N03

20 40 60 80 100 120 min. Time

Fig. 2-9 Results of Preliminary Model Computation (curves) vs. Palin's Data (symbols) at pH 8,5, Co/No=1.82

Cf = free chlorine, lVl = monochloramine, D = dichloramine, T = nitrogen trichloride, Ct = total chlorine

data that will serve best to eluc e the mechanism better.

As shown in Figures 2-5 through 2-9, the computed curves all

indicate a maximum in NHCl2 a few minutes after the beginning

of the reaction, followed hy a rather rapid decline. Palin's

for ten minutes of reac all are a when the

concentration of NHCl2 is changing quite rapidly so that any

slight time error in sampling or quenching the reaction would

have had a magni effect on the measured concentration.

He could not known this of course. Improvement

understanding of the breakpoint reaction definitely requires

better data on the concentration of NHCl2, especially during

the ti stages of on,

Fig. 2-6 shows a pattern of free

what unexpected, At pH 6.5 and a molar

orine which some-

of 1,82, the

rapid decline of free chlorine followed by a definite

resurgence, apparently due the regeneration of free

orine in React:lon (M-7). If this uncommon pattern can be

experimentally detected, it will undoubtedly serve as solid

evidence in support of some aspects of the proposed mechanism,

Also, iments with initial ammonia concentrations

as greatly different as ib1e seem to be needed, S e,

aceording to the mechanism, the formation of NHC12 is

sec whi its decomposi first-order, vari-

ations in initial ammonia cone should have important

ts on the pattern of the reaction, Demonstration of such

would serve significantly as confirmation of important

of proposed mechanism,

In view of all these conclusions, it seemed advisable

not to make further model computations with changed rate

parameter until addi

secured

ex per data had been

2-JO

3-l

CHAP'fER III

EXPERIMENTAL l\llf:rHODS

As the conclus thEl end of Chapttn· II indicated,

with much emphasis on the initial s

arc needed to define the rate

more precisely, e experiments should

provide data on the changes in concentration of all reac

, and produc to as a eck with

model computations as possible,

From results of the preliminary model computations,

it appeared that the changes in concentration of free

chlorine and dichloramine rather unique patterns

that their det should rece spec

tion conduct of experiments,

TheGe cons led to the of

following analytical methods for the k stud:ies

s tion of the most appropriate from them,

A u l' . E' -, t' , 1 re. 1.m1.nary . va~ua ":Lone

l, Ultraviolet Absorption Spectrophotometry. Abe on

spectrophotometry gives both quant:i.tat and qualitative

information on constituents in a sampltol, because

wavelengths of light at which absorption maxima occur are

charac of the absorbing species, and magnitudes

of these maxima are measures of their concentrations, Also,

in general, the advantages of spectrophotometric measurements

include the ability to make successive measurements on

3-2

undisturbed samples with considerable rapidity. 'I'his rapidity

can be applied to success spectral scans order to

erve changes the progress of the reaction, or

to the measurement a single wavelength in order to follow

changes

Since dichloramine shows an orption maximum about

295 o/ with a molar absorptivity of about 310, a attempt

was made to monitor changes of absorbance at 295 mr during

the course of' breakpoint reactions, It was hoped that

change of' absorbance at this wavelength could be us to

demonstrate the maximum in concentration of dichloramine

predicted by the model computations, S""'""'" experimental

breakpoint reactions were conducted near pH?, and the

ent uv s trum, 220 to 3'-W Irif-• was scanned successively seven to eight times in initial 30 minutes or so of

reaction, Generally speaking, the absorbance within the UV

decreased monotonically at all wave-lengths, 'rhe

absorbance

of a

295 rrp- was also monitored continuously by use

did not show any maximum during the

ini stages of' reaction.

Inability to utilize UV absorbancE' for moni ing the

changes in concentration of NHCL, "'

probably due to the

spectral overlap of' the abs on of other ch cal species

th(, reacting systems, For instance, hypochlorite ion has

an absorpt maximum at about 291 rry;- w.i th a molar absorp··

i ty of about JIJ.O. S e hypochlorite ion decreases rapidly

while dechloramine is formed during the initial stage of

reaction, the increase of absorbance at 295 ~as a result

of the increase of NHCl2 could very well be more than

compensated by the decrease of OCl-.

2, !Vlethods :for Differentiation of Chloramines, Because of

the :futility in utilizing UV absorption spectrophotometry,

attention was turned to other analytical methods :for

differentiating chloramines. Generally, such methods are

much more complicated and time-consuming than UV absorption

measurements, It was found, for example, that complete

differentiation of free chlorine, monochloramine, and

dichloramine by the chemical methods eventually chosen

usually required a sample volume of 100 ml, and took about

6 minutes or more for completion of the determinations,

'.rhis relatively long period of time requinJd for obtaining

a single set of data is a serious disadvantage when it is

intended to obtain as many data as possible during

initial stages of reaction,

'rhree such methods are compared in the :following

paragraphs.

a. Amperometric Titration Method, 'l'he amperometric

method is a s adaptation of the voltammetric principle,

In a serial set of titratlons, the residual chlorine

titrated with a standard solution of phenylarsine oxide,

and a special galvanic cell is used to detect each end-

point. Free available chlorine, or HOC1 + OCl-, is deter-

.3-.3

mined by titrat at pH between 6, 0 and ? , 5, a range

which chloramines do not react directly w.i .. th the ti

The fferenti.ation of mono- and di-chloramine depends on

the fact that monochloramine reacts more readily wi

does dichloramine, So monochloramine can be

also by phenylars on of' a

iodide

amount of' KI into the sme sample, If' the pH of' the sample

to .3, 5-4, 5 and more KI is

can be subsequently in a similar manner(1?),

Accurate determinations of free chlorine cannot

in the presence of NC1 3 , which ti.trates partly as free

The trodes and agitator are

during the course of titration; there

in the s

, the cl

their surfaces omes crucial every determination,

made

of

especially if numerous samples are to be success-

ivelY a relatively short period of time.

b, NO'l'-FAS Method, The NOT-FAS method(1?) was

by Pal and subsequently us his

s ( 6, ? ) , The edure involves three

sucess of titration with colorimetric indicator•

In t, free ch1

are titrated with standard

with or:t!}otolidine as

and a nronor-tcion of the NC1 3

8 on at pH 6,3 to 6.5

In the second step NH2Cl

after the addition of a l amount of iodide,

liberated dine reacting with the i.ne, The

3-4

third step requires, first, the addition of acid to permit

the NHC12 to react with the iodide; then, the adjustment

of the pH back to 6,5 :for titration.

A separate sample is required to estimate NCl3, Either

of two procedures may be utilized, The first involves the

extraction of the NCl3 with CCli+ and subsequent titration

of the water layer for free chlorine, The second depends

on the destruction of the freo rosidual chlorino with oxalic

acid, leaving NC1 3 for subsequont titration as if it woro

freo chlorine.

The determination of NHC1 2 in the third step is rather

complicated, involving two adjustments of pH and proper

waiting. Only part of the NCl3 is recovered in the free

chlorine step. At times the rest of the NC1 3 recovered

as NHCl2,

c, DPD-FAS Method, The DPD-FAS (18, 19) scheme is

essentially a modification of NOT-FAS, with diethyl-E.-

phenylenediamine (DPD) replacing neutral orthQtolidine as

the indicator.

The procedure also involves three successive steps of

titration: In the absence of iodido ion, free chlorine rtlac

instantly with DPD at pH 6.4 to produce a red color.

Titration with ferrous solution results in a disappearance

of the red color at the end-point. Subsequent addition of

a small amount of iodide acts catalytically to cause NH2Cl

to produce the red color. After titration to disappearance

3-5

of color again wit.h F ~ further addition of iodide to

excess evokes a rapid development of color equivalent to

NHC1 2 , which can then o be ti

Unlike the reaction with neutral orthoto