Bromine and Bromine-Releasing Biocides
Jonathan N. Howarth
Michael S. Harvey
Bromine disinfectant products have widespread commercial use. This chapter will be dedicated to those compounds that contain or release bromine in the +1 oxidation state. These are known as oxidizing bromine biocides. Important nonoxidizing bromine biocides, such 2,2-dibromo-3-nitrilopropionamide, and nonoxidizing bromine preservatives, such as 2-bromo-2-nitro-1,3-propanerdiol and β-bromo-β-nitrostyrene, have been discussed in chapter 17 of the fourth edition.1 Elemental bromine does not occur naturally, but its bromide salt precursors are distributed in trace quantities throughout the earth. Seawater, for example, has a bromide content of approximately 65 ppm. Note that this is much less than the chloride content of seawater at 20 000 ppm. Although bromine can be (and still is) produced from seawater, certain bodies of water and underground formations contain much higher concentrations of bromide and serve as the major source for the bromine and bromine-based products manufactured today. The Smackover brine formation in the south central United States (Arkansas) is a concentrated source of bromide ion (4000-5000 ppm). Other sources of bromide include deep wells in the Great Lakes region of the United States and the Dead Sea in the Middle East. The Dead Sea is a particularly rich source with concentrations ranging from 5000 to 6500 ppm.2 The bromide ion is further enriched using evaporative ponds. In the 1950s, bromine extraction commenced in both Arkansas (United States) and Israel to take advantage of these abundant sources of bromide. In the late 1990s, bromine extraction from the Dead Sea also commenced in Jordan.
The production of bromine requires oxidation of the bromide-containing brine. This can be accomplished by several methods. Commercial quantities of bromine were first produced in the United States in 1846 and in Germany in 1865 using a combination of manganese dioxide and sulfuric acid. In 1890, Herbert M. Dow built a small plant in Canton, OH, that produced bromine by an electrolytic process. Today, the only oxidant of any commercial importance employs chlorine gas fed countercurrent to the flow of brine in a bromine extraction tower.3 According to the Arkansas Geological Commission, US bromine production in 2001 was 212,000 metric tons, with Arkansas’ output accounting for 97% of US production and about 40% of that worldwide.4
Although bromine was discovered almost 200 years ago, it is only within the last 35 years that bromine-based biocide technologies have achieved significant commercial use in water treatment. Elemental bromine itself has not been employed commercially to any significant extent for control of microorganisms due to hazardous management issues (it is a corrosive, fuming liquid). The hazards associated with elemental bromine have not posed a detriment to innovation but rather have led to the development of a wide variety of bromine-based delivery chemistries which are safer, less hazardous, and more convenient to use than the elemental form. Today, the water treatment community has many bromine-based biocide technologies to choose from—liquid two-component systems such as activated sodium bromide (NaBr) (effectively, in-situ generation of bromine), solid hydantoin-based technologies, and single-feed, liquid products such as sulfamic acid-stabilized bromine. Bartholomew5 reviewed the use of bromine chemistry in cooling water systems and this was subsequently updated to reflect the current state of the technology.6 The review chronicles the development of bromine-based disinfectants from the earliest inception to the state of the art, up to 2004.
An update of the commercial developments of bromine biocides used in cooling water systems will be provided in this chapter, in addition to reviewing other applications where they are used. These include
recreational water (swimming pools and spas)
municipal wastewater treatment
food safety
Other relevant topics in this chapter include the analytical methods for measuring bromine residuals, bromine-containing disinfection by-products (DBPs), and an overview of the worldwide regulatory outlook for bromine biocides.
WATER TREATMENT—EARLY STUDIES AND FUNDAMENTAL PRINCIPLESAlthough elemental bromine itself finds little application in water treatment, its use was suggested as far back as 80 years ago. In 1935, Henderson7 patented a process for treating water with bromine “to destroy any pathogenic organisms that may be present.” Henderson data showed that the performance of just 0.25 to 0.5 ppm bromine with Escherichia coli-contaminated water was equivalent to 1.5 to 2.0 ppm chlorine. It was also claimed, with the amounts of bromine used (<5 ppm) in the water, that elemental bromine was expected to disappear due to the presence of organic matter present and any after-treatment was unnecessary. Concepts contained in the patent, improved effectiveness and rapid residual decay, still contribute to use of bromine chemistry today. Subsequent laboratory studies of bromine confirmed a broad range of activity over many types of microorganisms. Reports of its ability to deactivate E coli; to disinfect water; and to kill spore-forming bacteria, yeasts, and molds appeared in the 1930s and 1940s.8,9,10,11,12,13
The addition of bromine to water generates hypobromous acid (HOBr) and hydrobromic acid (HBr). Depending on the pH, HOBr can further convert to hypobromite (OBr–). In 1938, Shilov and Gladtchikova13 correctly measured a pKa value of 8.7 for the HOBr – OBr– conversion.
Br2 + H2O = HOBr + HBr
HOBr + OH– = OBr– + H2O pKa = 8.7
A few years earlier, others accurately determined the pKa for the analogous chlorine system as 7.5.14
Cl2 + H2O = HOCl + HCl
HOCl + OH– = OCl– + H2O pKa = 7.5
The relative amounts of the two hypohalous acids will therefore vary with the system pH. Above pH 7.5, the relative concentration of hypochlorous acid (HOCl) declines rapidly, although this decline does not occur until above pH 8.7 with HOBr. The importance of the differences in the acid dissociation constants will be discussed graphically in the section on cooling water treatment.
The impact of pH on the effectiveness of chlorine had actually been known for many years (see chapter 15). Workers as far back as 1921 noted that high pH decreased the microbiological activity of hypochlorite.15,16 A striking example of this is the work by Rudolph and Levine17 on spores of Bacillus metiens. Time to produce a 99% kill with 25 ppm available chlorine varied from 2.5 minutes at pH 6 to 131 minutes at pH 10. There was a dramatic change in performance from pH 8 to pH 9, a pH range at which the majority of the cooling water treatment programs run at today. The authors concluded that the rate of kill of chlorine was directly related to the concentration of HOCl, which decreases rapidly at elevated pH. Many other studies over the years have pointed to a reduction in performance with chlorine at high pH levels.18 Additional studies pointed to improved performance of chlorine and bromine mixtures in the presence of ammonia (NH3) and other nitrogenous-containing materials.19,20,21 The chemistry of chlorine and bromine differs significantly in the presence of excess NH3. Chlorine forms predominately monochloramine, which is a relatively ineff ective biocide—some 50 to 100 times less active than free chlorine.22 Bromine, in contrast, produces a mixture of bromamines in rapid equilibria (mainly mono- and dibroamamine at pH 7 to 9), which are relatively effective biocides. Dibromamine, for example, is said to have the same activity as HOBr itself.23 At pH 8.2 and a mole ratio of NH3 to diatomic bromine (Br2) of 10 (ie, 1.1 ppm NH3 to 1.0 ppm Br2), the mixture consists of a 50:50 mixture of mono- and dibromamine.24 Improved microbiological effectiveness of bromamines versus chloramines was demonstrated against an E coli strain at pH 8.2.25
Another feature of bromamines is that they typically decay much faster in the environment than chloramines.24 The following equations summarize the chemistry of the haloamines as formed from chlorine or bromine in excess NH3.
NH3 + HOCl = NH2Cl + H2O (fast reaction, chloramine decays slowly)
NH2Cl + HOCl = NHCl2 + H2O (slow reaction, chloramine decays slowly)
NH3 + HOBr = NH2Br + H2O (fast reaction, bromamine decays rapidly)
NH2Br + HOBr = NHBr2 + H2O (fast reaction, bromamine decays rapidly)
In the mid-1980s, environmental concerns caused the shift from acid feeds with chromate to acid feeds with chromate/zinc and finally to totally chromate-free, alkaline-based cooling water treatment programs.26 These alkaline-based programs relied on polyphosphates and, later, phosphonates and copolymers for corrosion and scale control.27 It was clear that the chlorine-based technologies were less effective at the new higher pH environment that was now typically around pH 8.5 to 8.8, compared to the acid-assisted pH 6.0 to 7.0 employed with chromate-based programs. Figure 17.1 is a graphical representation of the pH/CO2 solution equilibria.28
It explains why discontinuation of acid causes pH to naturally rise into this range. As the hot water return cascades down the cooling tower fill, dissolved CO2 is stripped out and escapes to the atmosphere along with the evaporated water. The CO2 solution equilibria shift and stabilize around pH 8.5 to 8.8. When this occurs, the cooling water contains principally bicarbonate (HCO3–) alkalinity.
It explains why discontinuation of acid causes pH to naturally rise into this range. As the hot water return cascades down the cooling tower fill, dissolved CO2 is stripped out and escapes to the atmosphere along with the evaporated water. The CO2 solution equilibria shift and stabilize around pH 8.5 to 8.8. When this occurs, the cooling water contains principally bicarbonate (HCO3–) alkalinity.
 FIGURE 17.1 Carbon dioxide (CO2) solution equilibria as a function of pH. Mole fraction of dissolved inorganic carbon (DIC). Abbreviations: H2CO3, carbonic acid; HCO3–, bicarbonate; CO32-, carbonate. |
Increased energy costs also spurred the search for newer, more cost-effective technologies. It was realized that if system surfaces could be maintained under cleaner conditions, energy savings could outweigh the increased costs of biocidal treatment. Older cooling tower designs used plastic fiberglass or wood to break the water into smaller droplets of high surface area:volume ratios to facilitate heat transfer. High-efficiency film fill systems replaced these older splash fill designs. Film fill systems maximized cooling tower performance but placed further demands on biocide programs.29 It was essential that the thin gap between adjacent films had to be kept free of slime that could cause clogging and reduce cooling tower efficiency. In extreme cases, when biocontrol is lost, a film fill tower could even collapse under the weight of slime that accumulated.
 FIGURE 17.2 Solid forms of the halogenated hydantoins, 1-bromo-3-chloro-5,5-dimethylhydantoin (20 g tablets, left) and 1,3-dibromo-5,5-dimethylhydantoin (granules, right). |
Water recycling measures, especially those related to the use of municipal waste water as cooling water makeup, prompted intense industrial research into biocides effective in NH3 and organic nitrogen rich environments. To accommodate the needs of this evolving landscape, it was clear that conventional chlorine programs could not meet these challenges. The spotlight fell front and center on bromine biocides. Due to volatility, handling difficulties, corrosivity, and expense, there had been little interest in elemental bromine as an industrial biocide up to then. Researchers focused on easy-to-handle, safer, nonvolatile, less corrosive, less expensive bromine-releasing biocidal products.
BROMINE PRODUCTS AND BROMINE-RELEASING PRODUCTSHalogenated Hydantoins
The first halogenated hydantoin introduced for water disinfection was 1-bromo-3-chloro-5,5-dimethylhydantoin (BCDMH). It was generally sold as 1″ 20 g tablets first for recreational water treatment and later for cooling water disinfection. Subsequently, 1,3-dibromo-5,5-dimethylhydantoin (DBDMH) was commercialized as the demand for product with a higher amount of free halogen was needed. Compared to BCDMH, the higher amount of bromine in DBDMH reduced its solubility and was therefore introduced as a high surface area granular product to ensure that sufficient bromine was released into the water being treated (Figure 17.2). A comparison of the properties of both hydantoins is provided in Table 17.1.
TABLE 17.1 Properties of halogenated hydantoins | ||||||||||||||||||||||||||||||
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Another halogenated hydantoin product based on 1,3-chloro-5-ethyl-5-methyllhydantoin is available in the form of briquettes.30 These require no binder in manufacturing. Due to its higher chlorine content, this product is more soluble (about 0.3%) than BCDMH alone. The 98% active product is formally registered as BCDMH (60.0%), 1,3-dichloro-5,5-dimethylhydantoin (27.4%), and 1,3-dichloro-5-ethyl-5-methylhydantoin (10.6%).
BCDMH is sparingly soluble in water. The chemistry of BCDMH can be described as a facile release of HOBr, followed by the sluggish oxidation of spent Br– ion by monochloro DMH (Figure 17.3).
HOCl + Br– = HOBr + Cl–
Analysis of BCDMH in demand-free waters by the free and total the N,N-diethyl phenylenediamine (DPD) methods indicated that about half of the oxidant value analyzes as free halogen (ie, free HOBr) with the other half analyzing as combined halogen (ie, combined chlorine as chloro-DMH). Excess bromide ion and elevated water temperatures promote the hydrolysis of chloro-DMH with subsequent generation of HOBr. Cooling towers having short Holding Time Index (HTI; residence times) waste chloro-DMH to the blowdown. Nevertheless, the combination of HOBr and combined chlorine chemistries are claimed to provide good penetration into and control of biofilms encountered in industrial water systems.31 On the other hand, swimming pools have a near infinite residence time, losing water only to physical means (leaks, filter backwashing, bather drag out). These conditions allow full utilization of the halogen atoms of BCDMH into HOBr.
 FIGURE 17.3 Mechanism of 1-bromo-3-chloro-5,5-dimethylhydantoin hydrolysis to hypohalous acids in water. |
DBDMH is sparingly soluble in water. The product contains more available bromine activity than bromine itself—111% for DBDMH versus 100% for elemental bromine. Only one of the bromine moieties in Br2 is manifest as HOBr in water, whereas DBDMH reacts with water to readily release two moles (Figure 17.4).
Virtually all of the halogen residual produced from DBDMH analyzes as being free, as determined by the DPD method, in contrast to other hydantoin products. In many applications, this property leads to lower relative product consumption than other hydantoin products and thus less product inventory and handling (Figure 17.5).
 FIGURE 17.4 Mechanism of 1,3-dibromo-5,5-dimethylhydantoin hydrolysis to hypobromous acid in water. |
 FIGURE 17.5 Relative distribution of free and total halogen residuals in demand-free water (using the N,N-diethyl phenylenediamine method). BCDMH, 1-bromo-3-chloro-5,5-dimethylhydantoin; DBDMH, 1,3-dibromo-5,5-dimethylhydantoin. |
The DBDMH has lower water solubility than the BCDMH. Although a proprietary method was developed for making DBDMH tablets of excellent strength and structural integrity,32 their low surface area meant insufficient bromine residual was released into the treated waters. Consequently, high surface area coarse granules were introduced to the market (see Figure 17.2). Many industrial water treaters elect to classify the microbiological quality of the treated water in terms of the amount of free-chlorine (DPD method) present. Because all the halogen in DBDMH reports as free-chlorine, whereas only half reports as free-chlorine for BCDMH, the former chemistry is often preferred, even though it is more expensive (because it contains twice the amount of bromine as BCDMH). Another property of DBDMH is the fact that saturated solutions are near-neutral in pH and low in halogen odor. This contrasts with other solid products available in the marketplace, which often yield acidic solutions with a strong irritating halogen odor.
TABLE 17.2 Properties of sodium bromide solutions | ||||||||||||||||||||
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Activated Sodium Bromide
The NaBr is a white, crystalline inorganic salt (Table 17.2). It is readily soluble in water and sold as a clear, pH-neutral solution containing 38 to 42 weight percent (wt%) product. Other sources of NaBr include NaOH-neutralized solutions of HBr vapors, which are the by-product of brominated flame-retardant manufacturing. The NaBr is not a biocide itself but must be used in conjunction with an activating agent such as chlorine gas, sodium hypochlorite (NaOCl) solutions (bleach), calcium hypochlorite, ozone (O3), etc. The result of this dual feed approach is the generation of HOBr. Chemical activation examples are discussed in the cooling water uses section in this chapter.
Sulfamic Acid-Stabilized Bromine Products
To avoid the complexity of the two-component activated NaBr-oxidant system, research effort focused on single-feed liquid bromine products. The first products were based on perbromide (Br3–) salts that quell the vapor pressure of elemental bromine.33,34 The products were wasteful in bromine because only one of the 3 bromine moieties in Br3– materializes as biocidal HOBr on hydrolysis in water. The highly acidic nature of these products further limited their appeal and the products were never US Environmental Protection Agency (EPA)-registered in the United States. STABR-EX was the first sulfamic acid-stabilized liquid bromine biocide that was commercialized; it was a breakthrough product of its era.35 STABR-EX was manufactured by NaOCl oxidation of NaBr followed by stabilizing the mixture with sodium sulfamate. Unlike the acidic perbromide formulations, it was of high pH, possessed no headspace bromine vapors, and all the bromine values were exhibited as oxidizing bromine. Figure 17.6 compares the appearance of an acidic perbromide with that of a sulfamic acid-stabilized product.
 FIGURE 17.6 Comparative appearance of an alkaline sulfamic acid-stabilized product (left) and early generation acidic perbromide formulation (right). For a color version of this art, please consult the eBook. |
There are four major sulfamic acid-stabilized bromine products commercially available. They are all single-feed liquids that require no external activation and were developed, in part, to overcome the complexities of dual feed activated NaBr, which requires the use of NaOCl or dichlorine (Cl2). As easily handled, preactivated liquids, these products are amenable to shock dosing, a feature not possible with the sparingly soluble halogenated hydantoins, which are generally administered through a by-pass feeder on a continuous or intermittent basis. The ease with which sulfamic acid-stabilized bromine products can be used, coupled with inexpensive shock dosing regimens, means they are rapidly gaining market share over solid bromine delivery systems. Sulfamic acid-stabilized bromine products are differentiated by their methods of manufacture and by their trade names: STABR-EX, STABROM 909, and BromMax 10.2 and 7.1 (Table 17.3). Compared to unstabilized NaOCl, these products typically have an excellent shelf-life greater than one year, even in hot climates (Figure 17.7).
TABLE 17.3 Properties of sulfamic acid-stabilized bromine products | |||||||||||||||||||||||||||||||||||||||||||||||||||||||
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As with NaOCl (or bleach) solutions, it is recommended to avoid contact of the concentrated materials with metals such as iron, copper, zinc, mild steel, and aluminum because these will be corroded over time. The transition metal ions released into solution could accelerate the deterioration of associated products. It is also recommended to avoid mixing with concentrated NaOCl solutions because this can give rise to exothermic decomposition. Finally, these products should be stored away from direct sunlight and other sources of heat to prevent degradation over time.
Sulfamic acid-stabilized bromine products can freeze at temperatures below 30°F (-1.1°C), so necessary precautions should be taken in cold climates. BromMax 10.2 is the exception. Because of the concentrated nature of this product, it crystallizes pure sodium N-bromosulfamate (Figure 17.8) above the freezing point of water. Therefore, this product should not be used if there is the possibility of exposure to cold weather. BromMax 10.2 is mostly sold into warmer climate countries. It contains
33% more active ingredient than other stabilized bromine products, and this higher activity has an economic advantage to contain overseas shipping costs. Far higher concentrations of bromine were possible through the use of trichloroisocyanuric acid (TCCA) as the oxidant.36,37 The TCCA is a solid (91% available Cl2) and does not dilute the product, as is the case for another stabilized bromine product where 12% NaOCl is used to affect the oxidation (see Table 17.3).
33% more active ingredient than other stabilized bromine products, and this higher activity has an economic advantage to contain overseas shipping costs. Far higher concentrations of bromine were possible through the use of trichloroisocyanuric acid (TCCA) as the oxidant.36,37 The TCCA is a solid (91% available Cl2) and does not dilute the product, as is the case for another stabilized bromine product where 12% NaOCl is used to affect the oxidation (see Table 17.3).
 FIGURE 17.7 Long-term storage of a sulfamic acid-stabilized bromine product versus a bleach (sodium hypochlorite [NaOCl]) solution (dichlorine basis) at 40°C. |
Subsequent to the commercialization of the four stabilized bromine products, another higher concentration (9.2% activity as Cl2) product was introduced for larger industrial cooling water uses. It has the tradename MAXXIS and is a mixture of sodium chloro- and bromo-sulfamates. MAXXIS also has the same low temperature limitation as BromMax 10.2.
 FIGURE 17.8 Hydrated crystals sodium N-bromosulfamate that can develop under cold conditions (6°C/45°F). For a color version of this art, please consult the eBook. |
 FIGURE 17.9 Trichloroisocyanuric acid/sodium bromide [NaBr] tablet and sodium dichlorisocyanurate/NaBr granules. |
Bromine-Releasing Isocyanurate Compositions38
Bromine-releasing isocyanurate compositions include TCCA or sodium dichlorisocyanurate (SDIC) containing compositions mixed with a source of bromine (NaBr) (Figures 17.9 and 17.10). Examples of these products are summarized in Table 17.4. When either bromine-releasing isocyanurate composition is added to water, they effectively generate HOBr through the reaction of HOCl (liberated from the chlorinated isocyanurate) and NaBr. Because both compositions contain an excess of chlorine, the treated water will always contain mixed halogens. At elevated pH (>7) and at high dilution, the regeneration of the Br– ion into HOBr by the excess chlorinated isocyanurate is sluggish (J.N.H., unpublished data, 1990). The granules are often found to be too soluble to be used in any type of automatic feeding device and are only amenable by manual broadcast, usually for shock dosing (see Figure 17.9). On the other hand, the different solubilities of NaBr and TCCA mean that the NaBr preferentially leaches out of tablets over time. This causes the tablets to disintegrate into a high surface area mound of product, which accelerates dissolution even further (J.N.H., unpublished data, 1990). Thus, it may not be possible for the certain tablet formulations to deliver a controlled, meaningful dose of halogen to the water.
 FIGURE 17.10 Structures of trichloroisocyanuric acid and sodium dichlorisocyanurate. |
TABLE 17.4 Properties of bromine-releasing isocyanurate compositions | |||||||||||||||||||||||||||
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Bromine Chloride
Typical properties of bromine chloride (BrCl) are listed in Table 17.5.39
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