romine 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
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 PRINCIPLES
Although 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
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–
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
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
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
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–
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.