|
Chloramines have long been used to provide a disinfecting residual in distribution
systems where it is difficult to maintain a free chlorine residual or where disinfection by-product
(DBP) formation is of concern. While chloramines are generally considered less reactive than
free chlorine, they are inherently unstable even in the absence of reactive substances. These
reactions, often referred to as “auto-decomposition”, always occur and hence define the
maximum stability of monochloramine in water. The effect of additional reactive material must
be measured relative to this basic continual loss process.
The net reaction, simplified from the approximately 14 individual reactions that govern it,
can be written as,
3 NH2Cl ? N2 + NH3 + 3 Cl- + 3 H+
In essence, auto-oxidation is nothing more than the "break point" reaction, occurring however, at
a much reduced rate. While relatively slow, it may still account for a substantial loss of residual
in distribution systems characterized by long residence times.
A thorough understanding of chloramine formation and auto-decomposition reactions is
important because this system is highly dynamic and especially susceptible to perturbations in
pH, ammonia, and chlorine concentrations. The complex set of reactions describing this govern
the initial formation of various forms of chloramine species and their transformation to other
forms with time, loss of chloramines by auto oxidation, and demand reactions as well as
formation of DBPs.
A kinetic model describing chloramine formation and auto-decomposition was developed
(Jafvert and Valentine, 1992; Vikesland et. al. 1998). This model (see table below) is based on studies
of isolated individual reactions and on observations of the reactive ammonia-chlorine system as a
whole. The model includes three principle reaction schemes as well as pertinent equilibrium
reactions:
| | 1. | Speciation reactions of HOCl with ammonia and the chlorinated derivatives of
ammonia, and the corresponding hydrolysis reactions. |
| | 2. | Disproportionate reactions of chloramine species and the corresponding back reactions. |
| | 3. | Redox reactions that occur in the absence of measurable free chlorine. |
Chloramine decay can be adequately predicted by the detailed mechanistic model.
However, its use is relatively complicated. Chloramine decay in the absence of demand
reactions, may, under most conditions, be estimated by a simple second order relationship. The
integrated form can be given in terms of a single coefficient ( kVCSC ) as :
The value of kVCSC ( Valentine Chloramine Stability Coefficient) increases with
decreasing pH and initial chloramine concentration that determines total free ammonia in the
system. It also increases with increasing total inorganic carbon and temperature. It can be
rationalized in terms of measurable parameters and known rate constants and therefore does not
need to be measured for each water.
Modeling chloramine decomposition kinetics
|
Reaction |
|
Rate Constant |
|
Reference |
| (1) |
HOCl + NH3 ===> NH2Cl + H2O |
|
k1= 1.5x1010 M-1 h-1 |
|
Morris and Isaac (1983) |
| (2) |
NH2Cl + H2O ===> HOCl + NH3 |
|
k2= 7.6x10-2 h-1 |
|
Morris and Isaac (1983) |
| (3) |
HOCl + NH2Cl ===> NHCl2 + H2O |
|
k3= 1.0x106 M-1 h-1 |
|
Margerum et al. (1978) |
| (4) |
NHCl2 + H2O ===> HOCl + NH2Cl |
|
k4= 2.3x10-3 h-1 |
|
Margerum et al. (1978) |
| (5) |
NH2Cl + NH2Cl ===> NHCl2 + NH3 |
|
kd* |
|
Jafvert (1985) |
| (6) |
NHCl2 + NH3 ===> NH2Cl + NH2Cl |
|
k6= 2.2x108 M-2 h-1 |
|
Hand and Margerum (1983) |
| (7) |
NHCl2 + H2O ===> I |
|
k7= 4.0x105 M-1 h-1 |
|
Jafvert and Valentine (1987) |
| (8) |
I + NHCl2 ===> HOCl + products |
|
k8= 1.0x108 M-1 h-1 |
|
Leao (1981) |
| (9) |
I + NH2Cl ===> products |
|
k9= 3.0x107 M-1 h-1 |
|
Leao (1981) |
| (10) |
NH2Cl + NHCl2 ===> products |
|
k10= 55.0 M-1 h-1 |
|
Leao (1981) |
| (11) |
HOCl <===> H+ + OCl- |
|
pKa= 7.5 |
|
Snoeyink and Jenkins (1980) |
| (12) |
NH4+ <===> NH3 + H+ |
|
pKa= 9.3 |
|
Snoeyink and Jenkins (1980) |
| (13) |
H2CO3 <===> HCO3- + H+ |
|
pKa= 6.3 |
|
Snoeyink and Jenkins (1980) |
| (14) |
HCO3- <===> CO3-2 + H+ |
|
pKa= 10.3 |
|
Snoeyink and Jenkins (1980) |
|
*kd = kH[H+] + kH2CO3[H2CO3 ] + kHCO3[HCO3-]
where kH2CO3 = 4x104 M-2h-1, kHCO3 = 800 M-2h-1
|
The Valentine Chloramine Stability Coefficient can be calculated for any water source to
define the limit the chloramine stability. Utilities may use this to calculate the effects of water
quality and chloramination practices (e.g. Cl/N ratio) on chloramine stability. Observed rates that
exceed those predicted kVCSC could also be used to point to problems in the system such as the
existence of oxidizable iron or organics. Observed rates of loss that are less than this may point
to the existence of relatively stable organic chloramines.
Source: A Guide for the Implementation and Use of Chloramines by Harms and Owen. © 2004 AwwaRF. Reproduced with permission. This material is presented solely for informational purposes. More details are available at http://www.awwarf.org/research/TopicsAndProjects/projectSnapshot.aspx?pn=2847. |
