Free chlorine, which is the chlorine present in water as hypochlorous acid (HOCl) or hypochlorite ion (OCl-), is an important disinfectant used in drinking water treatment. When chlorine gas (Cl2) is added to water, it rapidly hydrolyzes into hypochlorous acid and hypochlorite ion, which together constitute free chlorine. The relative proportions of hypochlorous acid and hypochlorite ion depend on the pH of the water. At typical drinking water pH of 6.5-8, both species coexist. Free chlorine provides a residual disinfectant that prevents recontamination during distribution in the water supply network. However, free chlorine also reacts with natural organic matter and bromide ion in source waters to produce various disinfection byproducts, including trihalomethanes and haloacetic acids, which are regulated due to potential health risks at high concentrations. Balancing efficient disinfection while minimizing regulated disinfection byproducts is an ongoing challenge for drinking water utilities. This article examines the chemistry of free chlorine, its use and importance as a drinking water disinfectant, reactions that remove free chlorine residual, and regulated disinfection byproducts that are formed.
Chemistry of Free Chlorine
When chlorine gas (Cl2) is added to water, it rapidly hydrolyzes to form hypochlorous acid (HOCl) and hydrochloric acid (HCl) according to the following reaction:
Cl2 + H2O → HOCl + HCl
Hypochlorous acid partially dissociates to form hypochlorite ion (OCl-) according to the following equilibrium:
HOCl ⇌ H+ + OCl-
The relative proportions of hypochlorous acid and hypochlorite ion depend on the pH of the water. At typical drinking water pH of 6.5-8, both species coexist. The sum of hypochlorous acid and hypochlorite ion concentrations represents the free chlorine residual.
Hypochlorous acid is a weak acid with pKa = 7.5 at 25°C. At pH 7.5, hypochlorous acid and hypochlorite ion exist in equal proportions. At lower pH, hypochlorous acid predominates, while at higher pH, hypochlorite ion predominates.
The disinfection power of free chlorine comes mostly from undissociated hypochlorous acid rather than hypochlorite ion. Hypochlorous acid can penetrate cell walls and react with vital amino acids in proteins, enzymes, and other cellular components to kill microorganisms. Hypochlorite ion is less effective as a disinfectant compared to the undissociated acid.
Free Chlorine Species Distribution
The relative distribution of hypochlorous acid and hypochlorite ion as a function of pH can be illustrated using the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
Where:
– pKa = 7.5 at 25°C
– [A-] = molar concentration of hypochlorite ion (OCl-)
– [HA] = molar concentration of hypochlorous acid (HOCl)
The percentage of hypochlorous acid decreases and the percentage of hypochlorite ion increases as the pH increases. At the pKa, the two species exist in equal proportions. The figure below illustrates the relative distribution:
| pH | % HOCl | % OCl- |
|---|---|---|
| 6 | 94% | 6% |
| 7 | 70% | 30% |
| 7.5 | 50% | 50% |
| 8 | 30% | 70% |
| 9 | 6% | 94% |
Free Chlorine Disinfection
Free chlorine is widely used as a disinfectant in drinking water treatment due to its potent germicidal properties and low cost compared to other disinfectants like ozone or UV irradiation. Free chlorine provides a residual disinfectant that prevents recontamination during distribution in the water supply network. It is effective for inactivating most viruses, bacteria, and protozoan parasites.
The effectiveness of chlorination depends on the chlorine dose, contact time, pH, temperature, and the nature of the microorganisms present. Because hypochlorous acid is a stronger disinfectant than hypochlorite ion, chlorination is more effective at lower pH. A typical target residual is 0.2-0.5 mg/L free chlorine in the distribution system. Higher chlorine doses and contact times are required when treating water with higher levels of contaminants that exert chlorine demand.
Chlorination is not equally effective against all microorganisms. Bacterial spores, protozoan oocysts like Cryptosporidium, and certain viruses are highly resistant to chlorination. Additional treatment barriers like filtration, flocculation, or UV disinfection are often needed to achieve targeted log reductions of these chlorine-resistant pathogens.
The Ct concept (concentration x contact time) is used to quantify disinfection efficacy. The Ct value needed to achieve a target log inactivation of a pathogen decreases as the concentration of disinfectant increases and the contact time increases. pH also affects Ct values, as hypochlorous acid is a more potent disinfectant than hypochlorite ion. Ct tables based on experimental data are used by utilities to determine required chlorine doses and contact times.
Advantages of Chlorination
Key advantages of chlorination for drinking water disinfection include:
- Effective inactivator of bacteria, viruses, and protozoa
- Provides residual disinfectant to prevent recontamination
- Easy to generate on site and simple to use
- Low capital and operating costs compared to ozone or UV
Disadvantages of Chlorination
Disadvantages of chlorination include:
- Not effective against some chlorine-resistant pathogens like Cryptosporidium
- Reacts with natural organic matter to form disinfection byproducts
- Forms toxic and corrosive disinfection byproducts
- Tastes and odors may develop from over-chlorination
- Storage and transport of gaseous chlorine poses safety risks if mishandled
Chlorine Reactions
When free chlorine is added to water, it begins reacting with various inorganic and organic species exerting a chlorine demand. Consumption of chlorine through these reactions decreases the free chlorine residual. The main reactions include:
- Reactions with reduced inorganic species like ferrous iron, manganese, nitrite, and sulfide
- Reactions with natural organic matter (NOM)
- Formation of regulated disinfection byproducts (DBPs)
Maintaining a sufficient free chlorine residual is important to deliver adequate disinfection. However, excessive chlorination can lead to issues like formation of DBPs, taste and odor problems, and corrosion of distribution system pipes. Careful chlorine dosing and monitoring of residual levels is needed to balance efficient disinfection and control of DBPs.
Inorganic Chlorine Reactions
Reduced inorganic species exert a rapid chlorine demand. Ferrous iron (Fe2+) and manganous manganese (Mn2+) are oxidized by chlorine to ferric (Fe3+) and manganic (Mn4+) states. One mole of chlorine oxidizes one mole of Mn2+ and two moles of Fe2+.
Nitrite (NO2-) reacts with hypochlorous acid to form nitrogen gas, nitrate, and chloride ion. This nitrification reaction consumes chlorine rapidly.
Sulfide (S2-) oxidizes to elemental sulfur, consuming one mole of chlorine per mole of sulfide:
S2- + 4HOCl → S + SO42- + 4Cl- + 4H+
If chlorine is overdosed relative to sulfide, further oxidation to sulfuric acid (H2SO4) can occur:
S + 6HOCl → H2SO4 + 6Cl- + 4H+
Natural Organic Matter Reactions
Natural organic matter (NOM) like decaying vegetation leaches into source waters and reacts with chlorine to form disinfection byproducts (DBPs). These reactions exert a chlorine demand and decrease the free chlorine residual over time. Higher NOM levels require higher chlorine doses to maintain a residual.
The main classes of DBPs formed from NOM are:
– Trihalomethanes (THMs) – chloroform, bromoform, bromodichloromethane, dibromochloromethane
– Haloacetic acids (HAAs) – dichloroacetic acid, trichloroacetic acid, monochloroacetic acid
– Other DBPs like cyanogen chloride and bromate (with ozone)
Regulated Disinfection Byproducts
While chlorine is an effective disinfectant for pathogens, it also reacts with natural organic matter and bromide ion present in source waters to produce various toxic byproducts. Chlorinated DBPs were first identified in drinking water in the 1970s.
The first DBPs regulated by the EPA under the Disinfectants and Disinfection Byproducts Rule were total trihalomethanes. The Stage 1 and Stage 2 DBPR lowered the MCL for total THMs to 80 μg/L and also regulated five HAAs at 60 μg/L.
Here are the regulated chlorination DBPs, their chemical structures, and maximum contaminant levels (MCLs):
| DBP | Chemical Structure | MCL (mg/L) |
|---|---|---|
| Chloroform | CHCl3 | 0.08 |
| Bromoform | CHBr3 | 0.08 |
| Bromodichloromethane | CHBrCl2 | 0.06 |
| Dibromochloromethane | CHBr2Cl | 0.06 |
| Dichloroacetic acid | C2H2Cl2O2 | 0.06 |
| Trichloroacetic acid | C2HCl3O2 | 0.3 |
| Monochloroacetic acid | C2H3ClO2 | 0.07 |
Many other unregulated DBPs form at lower concentrations. Brominated DBPs are more cytotoxic and genotoxic than their chlorinated analogs. Minimizing residence time in the distribution system and removing precursor compounds like NOM and bromide ion can help reduce DBP formation.
Control of Disinfection Byproducts
Various treatment techniques are used to control the formation of regulated disinfection byproducts while still maintaining sufficient disinfection. Strategies include:
- Precursor removal – coagulation/flocculation to remove NOM; anion exchange to remove bromide
- Disinfectant change or reduction – switch from chlorine to chloramines, ozone, or UV disinfection
- Operational changes – reduce residence time in distribution system; avoid excess chlorine dosing
- Blending low DBP source waters – mix surface and groundwater sources
NOM removal by enhanced coagulation/flocculation is one of the most effective strategies for controlling THMs and HAAs. Up to 60% removal of TOC can be achieved with optimized coagulant dosing, pH, and mixing.
Switching from free chlorine to monochloramine residual provides longer-lasting protection in the distribution system while reducing regulated THM and HAA formation by 50-75%. However, monochloramine is less powerful for inactivation of viruses and bacteria.
Alternative disinfectants like ozone and UV light minimize DBP formation but provide no residual disinfectant. They are often used with a low chloramine residual to provide distribution system protection.
Precursor Removal
NOM is the main precursor for THM and HAA formation. Enhanced coagulation/flocculation can remove 20-60% of TOC depending on source water conditions and coagulant dose. Higher coagulant doses are required at lower raw water pH. Ferric chloride is more effective than alum for organics removal.
Anion exchange resins can lower bromide concentrations in source water by over 90% to reduce formation of brominated THMs and HAAs. However, disposal of spent regenerant brine can be problematic.
Alternative Disinfectants
Switching to ozone or UV light followed by a chloramine residual can reduce THMs and HAAs by 60-80% compared to free chlorine. However, they provide no residual disinfection within the distribution system.
Chlorine dioxide (ClO2) forms lower levels of THMs and HAAs but produces inorganic byproducts like chlorite and chlorate. It cannot be used to maintain a distribution system residual.
Operational Strategies
Reducing water age in distribution system reservoirs and tanks decreases THM and HAA formation. Maintaining good turnover and possibly reconfiguring storage to eliminate stagnant zones can help.
Booster disinfection should be avoided unless required to maintain a minimum residual. Excessive chlorination can lead to taste and odor issues.
Conclusion
In summary, free chlorine is an essential disinfectant for inactivating pathogens in drinking water. However, chlorine also reacts with natural organic matter to produce various disinfection byproducts that are regulated due to toxicity concerns. Balancing efficient disinfection while controlling regulated and unregulated DBPs is an ongoing challenge for drinking water utilities. A variety of precursor removal, alternative disinfectant, and operational strategies can help reduce DBP formation while still providing adequate disinfection for safe drinking water. Continued research is needed to understand DBP formation mechanisms, toxicology, and effective control strategies.
