Archive for April, 2010

Breakpoint Chlorination Diagram

Break Point Chlorination

As chlorine is added, it is consumed by chemical reaction with the net effect of a rise in chlorine concentration. The slope will depend on addition rate and reaction rate. For the usual rates of addition, the reaction rate will suddenly speed up so that the concentration of chlorine falls. This is explained by the ease with which chlorinated compounds accept more chlorine. In other words, the rate of addition of the first atom of chlorine is relatively slow, but rates are greater for further reaction because chlorinating potentiates reactivity. When most of the reactions with chlorine are complete, the addition of more chlorine results in a permanent residual. A reasonable time for the experiment is 30 minutes. The point at which the concentration returns to an upward slope is termed the breakpoint. There may be no breakpoint observed for certain waters because different organic compounds react at various rates.

Breakpoint Clorination Disertation

Research Objectives
Overall Objectives of this study were:
 Determine the process dynamics for real-time, continuous-flow breakpoint operations.
 Define design criteria for a full-scale breakpoint facility.
 Predict VPDES permit compliance.
 Address safety issues.

Phase I
Specific Objectives for Phase I were:
 Determine optimum chlorine to ammonia dose ratio.
 Determine optimum reaction pH for the process
 Determine reaction times and required detention times.
 Investigate analytical methods for reliable measurement of parameters required for control
of the breakpoint process.
 Determine optimum S02:Cl2 dose ratio for dechlorination.
Phase II
Specific Objectives for Phase II were:
 Determine effects of different influent water temperatures (8, 12 and 20 °C) on the
performance of the breakpoint reaction.
 Determine the effect of various influent ammonia concentrations (2.0, 6.0 and 11.0 mg/L
NH3-N) on the breakpoint reaction.
 Determine the effect of increased influent nitrite concentrations (~5.0 mg/L NO2-N) on the
dose ratios required for breakpoint.
 Determine the effect of increased influent organic nitrogen concentrations (~1.0 mg/L) on the
breakpoint reaction.
 Perform a special study to assess the potential for nitrogen trichloride formation during
breakpoint operation.
Literature Review
There has been limited research done on the breakpoint chlorination phenomenon. Initial research
occurred in the first half of the century when breakpoint was a hot topic and drew the attention of
researchers. The latter half of this century has seen little done in the way of research into breakpoint
chlorination. With the exception of a few papers, mostly plant specific studies have been performed.
There has, however, been a great deal of research done in the field of chlorine and chlorine-ammonia
chemistry, as well as advances in the fields of chlorine and ammonia analysis. While this chapter will
certainly review all obtainable research on breakpoint chlorination, it will also focus on research done
in the field of chlorine-ammonia chemistry, as it pertains to the breakpoint reaction.
Breakpoint Chlorination
The physical-chemical process of ammonia oxidation with chlorine has been practiced in the water
treatment field for over 50 years. As early as the 1920s superchlorination was used as a successful
means of controlling taste and odors in water treatment plants. In the 1930s an unexplained
phenomenon was being observed at water treatment plants using higher than normal chlorine dosages.
These events prompted research into the chlorination reactions occurring at water treatment plants.
Among the first researchers to explain these chlorine reactions Griffin (16) used the term breakpoint
to describe the point where chlorine and ammonia concentrations were simultaneously minimized.
The breakpoint reaction is defined as the chlorination of a water containing ammonia resulting in an
initial increase in combined chlorine residual, followed by a decrease in the combined chlorine residual
along with ammonia concentrations, followed by an increase in free chlorine residual and near
complete removal of ammonia as nitrogen gas. Fig. 1 shows a hypothetical breakpoint curve for a
water with a dose requirement of 9:1 Cl : NH3 (20). Initial research efforts into the mechanism of the
breakpoint reaction are attributed to Calvert (4), and later studies by Griffin and Chamberland (17),
and Rossum (34). Ensuing research by others (32,41,43) has led to an understanding of the
stoichiometry and kinetics associated with the breakpoint process. More recently, a comprehensive
study of the kinetics of breakpoint chlorination was performed by Saunier and Selleck (36). The goal
of their work was to develop a mathematical model, derived from laboratory observations, which
would provide “a rational basis for the design and operation of the breakpoint process in order to
achieve predictable ammonia removal”(36). Unfortunately, past research yielded little or no insight
into the problem of successfully controlling the breakpoint process in a full-scale wastewater
treatment plant. Although Saunier and Selleck (36) performed comprehensive pilot study work, the
results were never incorporated into a full-scale plant application. Pressley et al. (33) performed
extensive pilot study research in order to provide design criteria for a full scale breakpoint operation
at the Blue Plains wastewater treatment plant in Washington, D.C.. Atkins et al. (1) performed an
extensive pre-design pilot study to provide information for full- scale breakpoint operations at the
Owosso wastewater treatment plant in Michigan. The engineering firm of Camp, Dresser & McKee
(5) also performed bench scale
Figure 1. Theoretical breakpoint curve. (Zone 1 is associated with the reactions of chlorine and
ammonia to form Monochloramine; Zone 2 is associated with an increase in dichloramine and the
disappearance of NH3; Zone 3 is associated with the appearance of free chlorine after the
breakpoint testing before construction of full scale facilities at the Lower Potomac Pollution Control
Plant in Lorton Virginia. However, no matter how successful the pilot studies (1,5,33) were, or how
much information they yielded, none have been able to use the pilot information to successfully
control a full scale breakpoint facility.
Chlorine Chemistry
Chlorine Hydrolysis
Sodium hypochlorite hydrolyzes rapidly in water according to the following reaction:
NaOCl + H2O ® HOCl + NaOH (1-1)
The formation of HOCl via the above reaction is essential before the initiation of the breakpoint
reaction. Because hypochlorous acid is a weak acid it undergoes only partial dissociation as follows:
HOCl º H+ + OCl– (1-2)
At pH values between 6.5 and 8.5, the above reaction is incomplete and both species are present to
some degree. The extent of the above reaction can be estimated from the following equilibrium
Dissociation of HOCl in water has been measured by several investigators and has been shown to be
temperature dependant. The following are hydrolysis constant data published by Morris (27):
Table 1 Temperature Dependency of pKa for HOCl
Temperature °

Physical – chemical conditions which promote the formation of NCl3 during breakpoint are: 1) low
initial system pH and 2) Cl2 : NH3 dose ratios above 10 – 12:1. Under normal breakpoint pilot
operations these conditions are not satisfied and hence, there may not be enough NCl3 present in the
system to measure. In light of these circumstances a special reactor tank (see Figure 10) was used
to create conditions which favor the formation of NCl3. Since NCl3 is mostly insoluble in water, and
it readily volatilizes under turbulent conditions, it is very difficult to detect the compound in that
medium. Since no calibration standards are available for NCl3 it is also difficult to measure
quantitatively. However, the literature suggested a means of qualitative and quantitative analysis of
nitrogen trichloride by means of sample extraction with carbon tetrachloride (CCl4) followed by
sample analysis using UV scanning spectrophotometry (9). The NCl3 was extracted from the head
space of the reactor tank via a PYREX® gas washing bottle filled with a known volume of
spectrophotometric grade CCl4. Once extracted into the CCl4, nitrogen trichloride is in a soluble form
and can be measured with a U.V. spectrophotometer, with a fingerprint of primary and secondary
peaks at 265 & 345 nm respectively. The literature also suggests that quantitative results can be
obtained through a ratio determination of primary and secondary peak absorbance (9).
Experimental Methods
1. The pilot reactor was set up using the existing final effluent tank at the pilot study. This
provided the mixer and access ports needed to conduct the experiment. The tank was sealed with
duct tape and removable rubber stoppers to provide access to add chemicals. A sampling port was
provided for a grab analysis to assure operating guidelines (pH, initial ammonia concentration and
Cl2:NH3 dose ratio) were met.
2. The water in the reactor tank (Fig.10) was UOSA final filter effluent of a known volume. A
pre-determined mass of ammonium chloride (NH4Cl) was added resulting in a final ammonia
concentration of ~11.0 mg/L NH3-N. Diluted sulfuric acid was added for pH adjustment. Calculated
volumes of 5.25% sodium hypochlorite were added to reach the breakpoint at the varying dose ratios
3. The reaction proceeded in the sealed tank for approximately 24 hours with constant, vigorous
mixing. During this time the head space of the tank was pumped, via a 12V DC Cole-Parmer Model
#7530-25 diaphragm pump and Tygon® tubing, through a gas scrubber filled with a known volume
of CCl4. This allowed any existing NCl3 in the system to transfer from the aqueous phase (in water)
to the gaseous phase in the head space and finally back to the soluble phase in CCl4. The gas exiting
from the scrubber was recirculated back into the reactor for reprocessing through the system. This
allowed maximum transfer of any existing NCl3 into the CCl4.
4. At the end of the 24 hour period the CCl4 was collected and analyzed using a scanning UV
spectrophotometer at wavelengths between 200 and 400 nm. The primary and secondary peaks of
NCl3 in CCl4 occur at 265 and 345 nm, respectively.
Results and Discussion
Batch testing was performed over a period of several days using several different Cl2:NH3 dose ratios
and different system pHs. While the results (Figure 11 and Table 4) show an increase in NCl3 as the
pH decreased and Cl2:NH3 dose ratio was increased, production of NCl3 appears to be more
dependant on low system pHs (<6.0) than higher dose ratios (>10:1). The most severe NCl3
formation appears when operating with both low system pH and a high Cl2:NH3 dose rate.
Operational data can be found in Appendix D.
UOSA currently uses NaOCl for disinfection and will use it for Breakpoint operations, low pH levels
are not a major concern. It is however, recommended that NaOH feed be available for situations
where high ammonia concentrations are encountered, to avoid any possibility of NCl3 generation.
Based on results from the batch testing it is recommended that the breakpoint facility be operated at
a high pH (>7.0 after 30 minutes detention) and low initial Cl2:NH3 dose ratios (<10.0, depending
on influent water quality) to prevent the formation of high concentrations of NCl3.

Phase I. Summary and Conclusions
Pilot Operations
· All pHs provided acceptable operation during steady state operation.
· Slower reaction rates were observed at lower pHs.
· Lower pHs were associated with less stable breakpoint operations.
· Alkalinity consumption was observed close to being stoichiometric.
· Chlorine addition in excess of the Cl2 demand plus the stoichiometric amounts required to reach
the breakpoint resulted in a corresponding decline in ammonia removal efficiencies.
· Data from pilot operation at higher pHs (7.5 & 8.0) showed :
§ lower 30 minute free and total Cl2 residuals,
§ lower effluent NH3 and TKN concentrations,
§ Cl2:NH3 closer to stoichiometric,
§ final effluent pH closer to neutrality (i.e. pH ~7.0).
Laboratory Data
· COD removal was observed through breakpoint.
· Minimal NO3
– production was observed during the breakpoint reaction.
· Total Solids increased through the breakpoint process.
· No Fecal Coliforms were detected after 30 minutes contact time in the pilot.
DAS Data
· Loss of breakpoint was observed more frequently at lower operating pHs.
· Loss of breakpoint at higher pHs was not as severe (easier to get back).
· The ability to run at lower free and total chlorine residuals was observed as operating pHs
increased from 7.0 to 8.0.
· Numerous unexplained perturbations occurred at pH 7.0, causing loss of breakpoint.
· A more stable breakpoint process was occurred at pH 8.0.
DBP Data
· Consistent increases in Chloroform through breakpoint, though not to levels that would
jeopardize drinking water sources.
· Creating Total Organic Halides through breakpoint, again, not to levels that would jeopardize
drinking water sources.

Water Aeration

Chapter Four - Aeration


Methods of Aeration
Aeration is a unit process in which air and water are brought into intimate contact. Turbulence increases the aeration of flowing streams (Figure 4-1). In industrial processes, water flow is usually directed countercurrent to atmospheric or forced-draft air flow. The contact time and the ratio of air to water must be sufficient for effective removal of the unwanted gas.

Aeration as a water treatment practice is used for the following operations:

  • carbon dioxide reduction (decarbonation)
  • oxidation of iron and manganese found in many well waters (oxidation tower)
  • ammonia and hydrogen sulfide reduction (stripping)

Aeration is also an effective method of bacteria control.


Two general methods may be used for the aeration of water. The most common in industrial use is the water-fall aerator. Through the use of spray nozzles, the water is broken up into small droplets or a thin film to enhance countercurrent air contact.

In the air diffusion method of aeration, air is diffused into a receiving vessel containing counter-current flowing water, creating very small air bubbles. This ensures good air-water contact for “scrubbing” of undesirable gases from the water.

Water-Fall Aerators

Many variations of the water-fall principle are used for this type of aeration. The simplest configuration employs a vertical riser that discharges water by free fall into a basin (Figure 4-2). The riser usually operates on the available head of water. The efficiency of aeration is improved as the fall distance is increased. Also, steps or shelves may be added to break up the fall and spread the water into thin sheets or films, which increases contact time and aeration efficiency.

Coke tray and wood or plastic slat water-fall aerators are relatively similar in design and have the advantage of small space requirements.

Coke tray aerators are widely used in iron and manganese oxidation because a catalytic effect is secured by contact of the iron/manganese-bearing water with fresh precipitates. These units consist of a series of coke-filled trays through which the water percolates, with additional aeration obtained during the free fall from one tray to the next.

Wood or plastic slat tray aerators are similar to small atmospheric cooling towers. The tray slats are staggered to break up the free fall of the water and create thin films before the water finally drops into the basin.

Forced draft water-fall aerators (see Figure 4-3) are used for many industrial water conditioning purposes. Horizontal wood or plastic slat trays, or towers filled with packing of various shapes and materials, are designed to maximize disruption of the falling water into small streams for greater air-water contact. Air is forced through the unit by a blower which produces uniform air distribution across the entire cross section, cross current or countercurrent to the fall of the water. Because of these features, forced draft aerators are more efficient for gas removal and require less space for a given capacity.

Air Diffusion Aerators

Air diffusion systems aerate by pumping air into water through perforated pipes, strainers, porous plates, or tubes. Aeration by diffusion is theoretically superior to water-fall aeration because a fine bubble of air rising through water is continually exposed to fresh liquid surfaces, providing maximum water surface per unit volume of air. Also, the velocity of bubbles ascending through the water is much lower than the velocity of free-falling drops of water, providing a longer contact time. Greatest efficiency is achieved when water flow is countercurrent to the rising air bubbles.


In industrial water conditioning, one of the major objectives of aeration is to remove carbon dioxide. Aeration is also used to oxidize soluble iron and manganese (found in many well waters) to insoluble precipitates. Aeration is often used to reduce the carbon dioxide liberated by a treatment process. For example, acid may be fed to the effluent of sodium zeolite softeners for boiler alkalinity control. Carbon dioxide is produced as a result of the acid treatment, and aeration is employed to rid the water of this corrosive gas. Similarly, when the effluents of hydrogen and sodium zeolite units are blended, the carbon dioxide formed is removed by aeration.

In the case of cold lime softening, carbon dioxide may be removed from the water before the water enters the equipment. When carbon dioxide removal is the only objective, economics usually favor removal of high concentrations of carbon dioxide by aeration rather than by chemical precipitation with lime.

Air stripping may be used to reduce concentrations of volatile organics, such as chloroform, as well as dissolved gases, such as hydrogen sulfide and ammonia. Air pollution standards must be considered when air stripping is used to reduce volatile organic compounds.

Iron and Manganese Removal

Iron and manganese in well waters occur as soluble ferrous and manganous bicarbonates. In the aeration process, the water is saturated with oxygen to promote the following reactions:

ferrous bicarbonate
ferric hydroxide
carbon dioxide

manganese bicarbonate

manganese dioxide

carbon dioxide


The oxidation products, ferric hydroxide and manganese dioxide, are insoluble. After aeration, they are removed by clarification or filtration.

Occasionally, strong chemical oxidants such as chlorine (Cl2) or potassium permanganate (KMnO4) may be used following aeration to ensure complete oxidation.

Dissolved Gas Reduction

Gases dissolved in water follow the principle that the solubility of a gas in a liquid (water) is directly proportional to the partial pressure of the gas above the liquid at equilibrium. This is known as Henry’s Law and may be expressed as follows:

Ctotal  =  kP


Ctotal     =  total concentration of the gas in solution

P        =  partial pressure of the gas above the solution

k         =  a proportionality constant known as Henry’s Law Constant

However, the gases frequently encountered in water treatment (with the exception of oxygen) do not behave in accordance with Henry’s Law because they ionize  when dissolved in water. For example:

«   »
carbon  dioxide
hydrogen  ion
bicarbonate ion
hydrogen  sulfide
hydrogen  ion
hydrosulfide ion
ammonium  ion
hydroxide ion

Carbon dioxide, hydrogen sulfide, and ammonia are soluble in water under certain conditions to the extent of 1,700,  3,900, and 531,000 ppm, respectively. Rarely are these concentrations encountered except in certain process condensates. In a normal atmosphere, the partial pressure of each of these gases is practically zero. Consequently, the establishment of a state of equilibrium between water and air by means of aeration results in saturation of the water with nitrogen and oxygen and nearly complete removal of other gases.

As the equations above show, ionization of the gases in water is a reversible reaction. The common ion effect may be used to obtain almost complete removal of these gases by aeration. If the concentration of one of the ions on the right side of the equation is increased, the reaction is driven to the left, forming the gas. In the case of carbon dioxide and hydrogen sulfide, hydrogen ion concentration may be increased by the addition of an acid. Bicarbonate and carbonate ions in the water will form carbon dioxide, which can be removed by aeration.

In a similar manner, an increase in hydroxyl ion concentration through the addition of caustic soda aids in the removal of ammonia.

Figures 4-4, 4-5, and 4-6 show the percentage of gas removal that may be obtained at various pH levels.

Gas removal by aeration is achieved as the level of gas in the water approaches equilibrium with the level of the gas in the surrounding atmosphere. The process is improved by an increase in temperature, aeration time, the volume of air in contact with the water, and the surface area of water exposed to the air. As previously indicated, pH is an important consideration. The efficiency of aeration is greater where the concentration of the gas to be removed is high in the water and low in the atmosphere.


Temperature significantly affects the efficiency of air stripping processes. Therefore, these processes may not be suitable for use in colder climates. Theoretically, at 68°F the carbon dioxide content of the water can be reduced to 0.5 ppm by aeration to equilibrium conditions. This is not always practical from an economic standpoint, and reduction of carbon dioxide to 10 ppm is normally considered satisfactory.

Although removal of free carbon dioxide increases the pH of the water and renders it less corrosive from this standpoint, aeration also results in the saturation of water with dissolved oxygen. This does not generally present a problem when original oxygen content is already high. However, in the case of a well water supply that is high in carbon dioxide but devoid of oxygen, aeration simply exchanges one corrosive gas for another.

The efficiency of aeration increases as the initial concentration of the gas to be removed increases above its equilibrium value. Therefore, with waters containing only a small amount of carbon dioxide, neutralization by alkali addition is usually more cost-effective.

The complete removal of hydrogen sulfide must be combined with pH reduction or chemical oxidation.

Nonvolatile organic compounds cannot be removed by air stripping. For example, phenols and creosols are unaffected by the aeration process alone.

Table of Contents
(Chapter 03 Applying Quality Methods) (Chapter 05 Clarification)

Is The Chlorination Of Amino Compounds Diffusion Controlled?

Is The Chlorination Of Amino Compounds Diffusion Controlled?

X.L. Armesto, M. Canle L., J.A. Santaballa

Departamento de Química Fundamental e Industrial. Facultade de Ciencias, Universidade da Coruña. A Zapateira, s/n. E-15071 A Coruña. ESPAÑA



The chemistry of N-halo compounds has scarcely received any attention since Challis and Butler [1] wrote in 1968 “There have been no detailed investigations of the mechanism of amine halogenation…”. This has been so even in view of their widely recognized carcinogenic and / or mutagenic properties[2,3]. These compounds are generated during water disinfection with halogens, and under different conditions give place to several other toxic substances through different pathways (disproportionation[4,5], Grob fragmentation[6] and inter- and intramolecular (1,2)-elimination reactions[7,8] have been identified).

The chlorination of amines, amino acids and peptides by HOCl (the most commonly used water-disinfecting agent) to yield the N-Cl-compounds has been shown [9,10,11] to proceed via direct transfer of Cl from the O of the HOCl to the unprotonated N of the N-compound.

Scheme 1.

Notwithstanding the few recent studies available, some uncertainties still remain about this mechanism, one of the most important being whether the reaction is difussion-controlled or not. Here, we present some evindences to discard the difussion-controlled process in favour of a chemically controlled one.


Reagents. Commertial amino compounds of the best quality available were used without further purification. Aqueous chlorine solutions were used as chlorinating agent. These were generated every 3 or 4 days by bubling Cl2 (g) through a NaOH solution and spectrophotometrically titrated daily (λmax(H2O)=292 nm, ε=350 dm3.mol-1.s-1 for ClO when pH>12)[12]. CH3COOH / CH3COO, H2PO4 / HPO4-2, H3BO3 / NaOH and HCO3/ CO3-2 buffers, as well as NaOH solutions were used to control the pH. The total buffer concentration was kept about 0.02 in all cases and the pH>5 to avoid the presence of Cl2 (aq). Ionic strength was kept constant to 0.5 using NaClO4.

Equipment. SF-61 single-mixing and SF-61MX multi-mixing stopped flow spectrophotometers from Hi-Tech Scientificreg. were used to follow the reaction. This was done either by measuring the increase in absorbance due to the formation of the N-Cl-compounds or the decrease due to the consumption of the chlorinating agent. The temperature was kept constant to within +/-0.05 K by means of water-flow thermostating devices. The pH measurements were carried out with previously calibrated combined glass electrodes. The achieved accuracy for pH was +/-0.02 units.

Working procedure. All reactions were carried out under second order conditions, i.e.: by simultaneous mixing of equal concentrations of both the chlorinating agent and the N-compound. The integrated second order rate equation was adequately fitted to the A/t data using the Marquardt non-linear optimization algorithm[13]. Each rate constant reported here is an average of, at least, five kinetic runs, the reproducibility being better than 4 %.


The chlorination of amino compounds is a second order reaction, one order relative to each of the reagents:

The rate constant changes dramatically with the acidity of the medium, passing through a maximum at a pH corresponding to the average of the pKa’s of both reagents (Figure 1).

Figure 1: profile of the reaction rate with the pH. The mechanism for this reaction has been previously discussed elsewhere9, concluding that the reaction proceeds as shown in Scheme 1 (vide supra). The theoretical expression deduced for the observed rate constant is:

where k is the bimolecular rate constant for the process and Ka and Kc are the ionization constants for the N-compound and for HOCl, respectively. By fitting the experimental data to this equation we have worked out the bimolecular rate constants for a wide variety of primary and secondary N-compounds, including amines, amino acids, peptides, chloramines and amides.

Figure 2: reactivity – basicity / nucleophilicity relationship in the chlorination of N-compounds.

A plot of [log (k/mol-1dm3s-1)] versus the pKa of the N-compound shows a clear relationship between the basicity / nucleophilicity of the N-compounds and their reactivity towards the chlorinating agent. This behaviour has been found in a range of 13 pK units for more than 70 compounds, as shown in Figure 2.

In view of this behaviour, it could be considered that when pKN>9.0, then:

and the reaction has reached the difussion-control limit. This hypothesis, that has been pointed out previously[14], could be reinforced by the values we have obtained for the activation enthalpies of this reaction once accounted for the acid-base equilibria of both reagents (Table 1).

Table 1: Activation parameters for the chlorination of different N-compounds by HOCl.
N-Compound ΔH / kJ.mol-1 ΔS / J.mol-1.K-1
Me2NH 15 +/- 5 -46 +/- 6
Me,EtNH 23 +/- 3 -18 +/- 2
Et2NH 26 +/- 6 -11 +/- 1
Pr2NH 29 +/- 2 -2 +/- 0.2
iPr2NH 30 +/- 4 -5 +/- 1
iBu2NH 26 +/- 6 -18 +/- 2
Iminodipropionitrile 10 +/- 1 -108 +/- 2
Gly.Gly 0 -101 +/- 20
Gly.Gly.OEt 0 -131 +/- 15

I.e.: the very small ΔH could correspond to a difussion-controlled process, although ΔS lies in a wider range, even for compounds that have very similar pKN values and for which pKN > 9.0.

Notwithstanding these facts, there seems to be enough evidence to discard a difussion-controlled mechanism:

a) The higher rate constants we have found are around 108-109 mol-1.dm3.s-1, still well below the generally accepted limit of difussion-control. Table 2 summarizes the rate constants for the chlorination of Me2NH by different chlorinating agents, showing that there is no apparent relationship between the reaction rate and the size of the chlorinating agent, as should occur if the reaction were difusion-controlled.

Table 2: Rate constants for the chlorination of Me2NH by different reagents.
Chlorinating agent k /mol-1.dm3.s-1
(N-Cl),(N-Me)-Benzenesulfonamide 0.0245.107
(N-Cl)-Succinimide 1.53.107
(N-Cl)-Quinuclidinium 2.019.107
HOCl 7.501.107
NH2Cl 80.107
Cl2 160.107

b) The small values of activation entalpy could be an due to solvation phenomena or to the participation of the solvent in the transition state. In the case of primary and secondary N-compounds a transition state like:

is very likely to take place and would be thermodinamically favoured. The molecules under study are quite polar and we have no way to account for the influence of the changes in solvation upon the activation parameters.

c) We have found a clear steric effect within a family of similar compounds, the reaction rate decreasing as the steric hindrance increases. Figure 3 shows this effect using Charton’s ν parameter. The effect is found both for the groups directly attached to the N (case of the amines) and for those attached to a C in the vicinity of the N (case of the amino acids), being more important in the first case. This behaviour does not agree with which would be expected for a physically-controlled reaction, but with a chemically-controlled one.

Figure 3: Steric effects in the chlorination of N-compounds by HOCl.

d) We have found values of (kH/kD)<1 for the solvent isotope effect in the chlorination of several N-compounds in the pKN range 3.90<pKN<12.4 by HOCl. Thus, for example, (kH/kD)=0.70 for Glycine (pKN=9.78). Again, this behaviour would not be expected for a difussion-controlled process.


Taking into account all the previous evidences, we come to the conclusion that the chlorination reaction through transfer of Cl from the O of HOCl to the N of the corresponding N-compound is a chemically controlled process, at least in the wide range of pKN we have studied.

Work in progress.

In view of the very special reactivity – basicity / nucleophilicity pattern observed, we are re-visiting this mechanism in terms of a halogen transfer reaction, as compared with other well-known group transfers. We are also carrying out more research in order to characterize the transition state.


[1] B. Challis, A.R. Butler. “Substitution at an amino nitrogen”, in “The Chemistry of the Amino Group”. Edited by S. Patai. Interscience Publishers. London, 1968.

[2] J. Owusu-Yaw, W.B. Wheeler, C.I. Wei. Water Chlorination. Environmental Impact and Health Effects. Vol. 6, pp. 179-191. Lewis Publishers. Inc. Chelsea, 1990.

[3] A.C. Sen, J. Owusu-Yaw, W.B. Wheeler, C.I. Wei. J. Food. Sci., 54 (4), 1057-1060&1065 (1989).

[4] J.M. Antelo, F. Arce, M. Parajó, P. Rodríguez, A. Varela. Int. J. Chem. Kinet., 24, 991-997 (1992).

[5] X.L. Armesto, M. Canle L., M. Losada, J.A. Santaballa. Int. J. Chem. Kinet., 25, 331-339 (1993).

[6] X.L. Armesto, M. Canle L., M. Losada, J.A. Santaballa. J. Org. Chem., 59, 4659-4664 (1994).

[7] X.L. Armesto, M. Canle, M. Losada, J.A. Santaballa. J. Chem. Soc., Perkin Trans. 2, 181-185 (1993).

[8] X.L. Armesto, M. Canle L., P. Carretero, M. Losada, J.A. Santaballa. Unpublished work.

[9] X.L. Armesto, M. Canle L., J.A. Santaballa. Tetrahedron, 49, 275-284 (1993).

[10] X.L. Armesto, M. Canle L., M.V. García, M. Losada, J.A. Santaballa. Int. J. Chem. Kinet., 26, 1135-1141 (1994).

[11] X.L. Armesto, M. Canle L., M.V. García, M. Losada, J.A. Santaballa. Gazz. Chim. Ital., 124, 519-523 (1994).

[12] J.C. Morris. J. Phys. Chem., 70, 3798 (1966).

[13] D.W. Mardquardt. J. Soc. Ind. Math., 11, 431 (1963).

[14] D. Matte, B. Solastiouk, A. Merlin, X. Deglise. Can. J. Chem., 67, 786-791 (1989).

Reactions of chlorine with inorganic and organic compounds

Reactions of chlorine with inorganic and organic compounds during water treatment-Kinetics and mechanisms: a critical review.

Marie Deborde, Urs von Gunten

Water Research (2008)

Volume: 42, Issue: 1-2, Pages: 13-51

  • PubMed ID: 17915284


Numerous inorganic and organic micropollutants can undergo reactions with chlorine. However, for certain compounds, the expected chlorine reactivity is low and only small modifications in the parent compound’s structure are expected under typical water treatment conditions. To better understand/predict chlorine reactions with micropollutants, the kinetic and mechanistic information on chlorine reactivity available in literature was critically reviewed. For most micropollutants, HOCl is the major reactive chlorine species during chlorination processes. In the case of inorganic compounds, a fast reaction of ammonia, halides (Br(-) and I(-)), SO(3)(2-), CN(-), NO(2)(-), As(III) and Fe(II) with HOCl is reported (10(3)-10(9)M(-1)s(-1)) whereas low chlorine reaction rates with Mn(II) were shown in homogeneous systems. Chlorine reactivity usually results from an initial electrophilic attack of HOCl on inorganic compounds. In the case of organic compounds, second-order rate constants for chlorination vary over 10 orders of magnitude (i.e. <0.1-10(9)M(-1)s(-1)). Oxidation, addition and electrophilic substitution reactions with organic compounds are possible pathways. However, from a kinetic point of view, usually only electrophilic attack is significant. Chlorine reactivity limited to particular sites (mainly amines, reduced sulfur moieties or activated aromatic systems) is commonly observed during chlorination processes and small modifications in the parent compound’s structure are expected for the primary attack. Linear structure-activity relationships can be used to make predictions/estimates of the reactivity of functional groups based on structural analogy. Furthermore, comparison of chlorine to ozone reactivity towards aromatic compounds (electrophilic attack) shows a good correlation, with chlorine rate constants being about four orders of magnitude smaller than those for ozone.

Industry News: National Swimming Pool Foundation Awards Six Research Grants Totaling $415,282

Industry News: National Swimming Pool Foundation Awards Six Research Grants Totaling $415,282 — March 24, 2009

COLORADO SPRINGS, Colorado, March 24. RECENTLY, despite difficult economic times, the National Swimming Pool Foundation (NSPF) board of directors awarded six grants totaling $415,282 to continue efforts to study the health benefits unique to aquatic exercise and immersion in hot/warm water, and to reduce injury and disease in and around the water. Research results will be reported by grant recipients at the 2009 World Aquatic Health Conference (WAHC) October 28-30 in Atlanta, Georgia. “Research helps spur long-term growth. Providing these grants is our version of an economic stimulus package – without taking on debt,” remarks Bill Kent, Chairman NSPF Grant Review Committee.

Four health benefit grants worth $350,282 were awarded to Utah State University, University of South Carolina, West Virginia University, and Washington State University – National Aquatic and Sports Medicine Institute. Two injury prevention grants worth $65,000 were awarded to Purdue University and University of North Carolina – Charlotte. These grants will help to sustain ongoing research supported by NSPF in recent years.

Bruce Becker, M.D., Washington State University, National Aquatic and Sports Medicine Institute (NASMI), was awarded a grant of $200,000 based on the NSPF 5-year commitment to help establish a world-renowned health benefit research center. This grant will support efforts to continue to understand hot water immersion, and also study the aquatic exercise effects on subjects with asthma. Dr. Becker reported on his second year hot tub immersion study findings at the 2008 WAHC, which can be viewed online at

Stephen N. Blair, P.E.D., University of South Carolina, was awarded a grant of $91,323. Dr. Blair reported on the first year results of the health benefits of swimming and mortality at the 2008 WAHC with compelling findings relative to men and lower all-cause mortality risk. This new grant will continue the research and study effects of swimming on disease and injury. Dr. Blair’s research analyzes data from over 75,000 individuals completed from the Aerobics Center Longitudinal Study, following them over a span of more than 32 years.

William Hornsby, Ph.D., West Virginia University, was awarded a grant of $38,824 to continue his study of the psychological and physiological effects of land versus water-based exercise with patients who have Type-2 Diabetes Mellitus. Exercise and diet are the key recommendations on the control and prevention of Type-2 Diabetes. Almost all studies focus on land-based exercise programs. Since most diabetes patients are obese, water exercise is ideal to reduce joint stress. Dr. Hornsby’s report of first year results was given at the 2008 WAHC and the video is available at

Dennis Dolny, Ph.D. and Eadric Bressel, Ph.D., Utah State University, were awarded a grant of $36,458 to examine the acute biomechanical, physiological, and psychological responses of people with osteoarthritis (OA). The study will use matched controls while walking on land, compared to walking in chest deep water on an underwater treadmill. This is the first study to compare subjects with and without OA walking at various speeds on land and on an underwater treadmill. With the fundamental solutions, the desire is to establish guidelines for creating a template for OA individuals to use when they participate in water walking programs. This study is an extension of Dr. Dolny’s work previously presented at the 2008 WAHC which can be viewed at

James Amburgey, Ph.D., University of North Carolina-Charlotte (UNCC), was awarded a grant of $45,000 to continue his work on filtration research to remove crypto from pool water and reduce the risk of disease transmission. In addition, NSPF will manage and administer an industrial research grant, which may exceed $200,000, raised through industry donations to pursue this next phase. The goal of the research is to give manufacturers, regulators, operators, and pool designers a better understanding of the capabilities of multiple types of filters to remove oocysts in the swimming pool environment with or without clarifier addition. They will be able to develop better product label instructions, standard operating procedures, and remediation strategies to reduce the risk of waterborne disease outbreaks, thereby protecting public health and increasing participation in aquatic activities. Interested parties may still make donations to NSPF to fund this research.

Ernest Blatchley III, Ph.D., Purdue University, West Lafayette, Indiana, was awarded a grant of $20,000 to extend his research on the treatment methods for removal of volatile disinfection by-products from swimming pool water. Researchers look to address some of the important knowledge gaps that exist relative to UV and chlorine applications in recreational water. The research focuses on chemical and photochemical reactions that form and destroy disinfection by-products. NSPF contributed matching funding in 2006 so that a portable device (mass spectrometer) could be purchased. In 2007/2008, the NSPF grant funded the testing of air in indoor aquatic facilities to verify the laboratory experiments. Work continues to help reduce exposure to hazardous chemicals and to improve indoor air quality

Over the past 5 years, the board of directors of the National Swimming Pool Foundation has given back over 3.5 million dollars to fund research. “Our investment is creating growth opportunities to the watchful industry members,” says Thomas M. Lachocki, Ph.D., CEO of NSPF. “Organizations who find ways to reduce illness, injury, chemical exposure and drowning are prospering in these difficult times.” Lachocki observes, “They are also creating a safer industry that will be positioned to prosper as our economy rebounds. Research results are revealing the path forward to those who are paying attention.”

About NSPF
The National Swimming Pool Foundation is a non-profit organization dedicated to improving public health worldwide and is the leading educator of aquatic facility operators and the chief philanthropic research sponsor in the aquatics field. NSPF works towards its mission to encourage healthier living through aquatic education and research with its collection of multi-lingual educational products and training, and sponsors the annual World Aquatic Health Conference. Visit and


Clearing the Air: Chloramine Control for Indoor Swimming Pools

By Tom Griffiths, Ed.D.
Director of Aquatics, Penn State University
President, Aquatic Safety Research Group, LLC

Perhaps the most perplexing and controversial problem facing heavily used indoor pools today is chloramine production. Chloramines cause obnoxious odors as well as skin, eye, and respiratory irritation that many swimmers mistakenly attribute to chlorine itself. When chloramine levels become troublesome (0.3-0.5 parts per million (ppm)), people begin to complain. And while much finger pointing takes place, often little progress is made in correcting the problem. Swim coaches and competitive swimmers blame the pool operators, while pool operators in turn blame the swimmers; pool chemists blame the ventilation systems, whereas those in charge of air handling blame the water chemistry.

But who truly is to blame? And how is the problem fixed? In reality, everyone is to blame when chloramines are produced, and everyone has a role in controlling them. While a heavily used pool may never be completely chloramine free, you can greatly reduce chloramine production through good pool management practices.

Chronic chloramines and the associated smell and irritation are caused by a variety of factors. Despite what many swimmers assume, the major cause of these problems is too little free chlorine rather than too much! “Free” chlorine, used to kill germs and help prevent the spread of waterborne illnesses, also oxidizes natural waste products from swimmers, including sweat, body oil, urine and other ammonia-nitrogen compounds. If the free chlorine levels are not sufficiently high to oxidize these nitrogenous wastes, the free chlorine combines with them to form noxious cholarmine compounds. Whenever someone calls me with a chloramine problem, the first thing I tell him or her is that once they shock their pool (shock treatment will be discussed below) they should maintain a free residual of 0.5ppm higher than usual. This higher level of chlorine usually does the trick.

Another remedy that is rarely used but very effective is to enforce soap showers prior to swimming. A soap shower will remove excess body oils and sweat, thus greatly reducing the amount of body waste going into the pool. Some pool chemists claim that if everyone showered prior to swimming, it would reduce the chlorine demand by 50%. So perhaps you could get your swimmers to shower if you told them to shower, not because they are dirty but rather because their body oils react with chlorine to produce the smell they hate. Along those lines, competitive swimmers produce a great deal of sweat when they train rigorously. There is absolutely no way to avoid this, so you must plan on combating the perspiration that they will normally and regularly produce.

Even with the cleanest swimmers and the best water chemistry, though, chloramines can be a problem. If you have an energy efficient air handling system that re-circulates the air, often the chloramines are re-circulated and trapped in the building because they cannot escape. Air handling systems must bring in lots of fresh air and exhaust full blast when the pool is busy. If this is not done, chloramines will keep building. If the air handling system does not significantly exceed existing ASHRAE standards, then a heavily used pool will probably have an air quality problem.

Ever notice how you don’t notice chloramine odors at an outdoor pool? As they say, “No harm, no foul.”

Once you have an abundance of chloramines, they are not easy to get rid off. Just like algae growth in swimming pools, the key is prevention. And just like “Layers of Protection” for drowning prevention we also need “Layers of Protection” against chloramines. To help prevent and rid your facility of chloramines once they develop, you may want to experiment with a combination of the following:

  • “Shock” more often with free chlorine. Shock treatment involves raising the free chlorine level to at least 10 times higher than the combined chlorine level. Weekly is best for most pools but it may be required even more often for extremely heavily used pools.
  • Use a non-chlorine shocking agent like the monopersulfate-based oxidizers. These reduce chloramines without adding chlorine. Many pool operators find alternating between traditional chlorine and the non-chlorine shocking agents works best.
  • Add volcanic ash to your sand filters. This holds the ammonia in the filter tanks rather than in the swimming pools. Zeolite works well but must be regenerated to be effective in the long run.
  • Granulated Activated Carbon (GAC) filters may also be added to your existing filtration/circulation systems to remove ammonia that produces chloramines in the pool water.
  • Anticipate heavy bather loads. When you know your swimming pool is going to be inundated with swimmers either by way of a swim meet, swimming lessons or a huge rental or party, take preventive action prior to the swimmers arrival.
  • Shock the pool and keep the free chlorine levels up extra high before the swimmers enter the pool. Insist that the group shower before entering. These preventive measures will do wonders in keeping chloramines formation to a minimum.
  • If you have a good clean source of fresh water, give your filters and extended backwash so that you drain off lots of water (up to 1/3 of your pool volume) and replace it with fresh water.
  • Vacuuming and brushing your pool daily also removes much of the dirt chlorine reacts with that your filters have not caught yet.

A word of caution — Many water companies are using chloramines to disinfect the water they supply their customers. If your source water is disinfected with chloramines, as many are, you have an uphill battle on your hands. The facility may need to strip the chloramines with a GAC filter as the water enters the building and before it enters the swimming pool.

Finally, most of the “ideal” ranges recommended for chlorine in public swimming pools are simply too low and just plain wrong. Heavily used pools often need 3.0 – 4.0 PPM in order to prevent chloramines. Try running yours higher for six months. You’ll be glad you did.

For more information, refer to Chapter 10: Superchlorination The Complete Swimming Pool Reference, Sagamore Publishing and