3.1 Chlorine Chemistry: Hypochlorous Acid, Free/Combined Residual & Breakpoint Chlorination

Chlorine Chemistry: Hypochlorous Acid, Free/Combined Residual & Breakpoint Chlorination

Disinfection is the selective destruction or inactivation of pathogenic microorganisms to prevent the transmission of waterborne disease. In drinking water and wastewater treatment across Colorado, chlorine remains the predominant chemical disinfectant due to its high oxidation potential, proven germicidal efficacy, and ability to maintain a protective residual throughout distribution networks.

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1. Aqueous Chlorine Reactions

When chlorine is introduced to water, several distinct chemical reactions occur depending on the specific chemical form applied—gaseous elemental chlorine ($\text{Cl}_2$), liquid sodium hypochlorite ($\text{NaOCl}$), or solid calcium hypochlorite ($\text{Ca(OCl)}_2$).

Gas Chlorine Hydrolysis

Elemental chlorine gas hydrolyzes virtually instantaneously (within tenths of a second) when dissolved in water according to the following reversible reaction:

$\text{Cl}_2 + \text{H}_2\text{O} \rightleftharpoons \text{HOCl} + \text{H}^+ + \text{Cl}^-$

This reaction produces hypochlorous acid ($\text{HOCl}$) and hydrochloric acid ($\text{HCl}$). Because hydrochloric acid is a strong acid that fully dissociates, feeding chlorine gas consumes natural alkalinity and lowers finished water pH. Every pound of chlorine gas added theoretically consumes approximately 1.13 to 1.4 pounds of alkalinity as calcium carbonate ($\text{CaCO}_3$).

Hypochlorite Salt Hydrolysis

Hypochlorite salts also generate hypochlorous acid upon contact with water, but their reaction dynamics affect pH differently:

• Sodium Hypochlorite (Liquid Bleach, 12.5%): $\text{NaOCl} + \text{H}_2\text{O} \rightleftharpoons \text{HOCl} + \text{Na}^+ + \text{OH}^-$
• Calcium Hypochlorite (High-Test Hypochlorite / HTH, 65%): $\text{Ca(OCl)}_2 + 2\text{H}_2\text{O} \rightleftharpoons 2\text{HOCl} + \text{Ca}^{2+} + 2\text{OH}^-$

Both hypochlorite salts produce hydroxide ions ($\text{OH}^-$), which raises the solution pH. Consequently, treating acidic or poorly buffered waters with hypochlorites can cause the process pH to rise, requiring pH adjustment to ensure optimal disinfection kinetics.

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2. Hypochlorous Acid Dissociation and pH Dependence

Hypochlorous acid ($\text{HOCl}$) is a weak acid that partially dissociates in aqueous solution to yield a hydrogen ion ($\text{H}^+$) and a hypochlorite ion ($\text{OCl}^-$):

$\text{HOCl} \rightleftharpoons \text{H}^+ + \text{OCl}^-$

The relative distribution of $\text{HOCl}$ versus $\text{OCl}^-$ is strictly governed by water pH and water temperature. The dissociation constant ($pK_a$) for hypochlorous acid is approximately 7.54 at 20°C (and 7.60 at 15°C).

pH 6.0:  [========================================] 96.5% HOCl   | [=] 3.5% OCl-
pH 7.0:  [=============================] 77.5% HOCl             | [========] 22.5% OCl-
pH 7.5:  [====================] 50.0% HOCl                      | [====================] 50.0% OCl-
pH 8.0:  [========] 21.5% HOCl                                  | [=============================] 78.5% OCl-
pH 8.5:  [===] 8.7% HOCl                                        | [==================================] 91.3% OCl-
pH 9.0:  [=] 2.7% HOCl                                          | [======================================] 97.3% OCl-

Why HOCl is 80 to 100 Times More Potent Than OCl⁻

Bacterial cell membranes carry a net negative surface charge. Because hypochlorous acid ($\text{HOCl}$) is an electrically neutral molecule, it effortlessly diffuses through the hydrophobic lipid bilayer of microbial cell walls. Once inside the cell, $\text{HOCl}$ oxidizes essential intracellular sulfhydryl-containing enzymes, disrupting glucose metabolism and cellular respiration. In contrast, the hypochlorite ion ($\text{OCl}^-$) carries a negative electrical charge and is electrostatically repelled by the microbial membrane, rendering it 80 to 100 times less effective as a germicidal agent than $\text{HOCl}$.

3. Chlorine Reactions with Ammonia: Chloramines

When water contains dissolved ammonia nitrogen ($\text{NH}_3\text{-N}$), hypochlorous acid reacts in sequential substitution steps to form inorganic chloramines, which constitute combined available chlorine residual.

Monochloramine Formation

At neutral to slightly alkaline pH (7.0 to 8.5) and a weight ratio of chlorine to ammonia nitrogen of up to 5:1, hypochlorous acid reacts rapidly with ammonia to form monochloramine:

$\text{NH}_3 + \text{HOCl} \longrightarrow \text{NH}_2\text{Cl} \text{ (Monochloramine)} + \text{H}_2\text{O}$

Monochloramine is a stable, mild disinfectant that does not produce objectionable chlorinous tastes or odors when maintained within proper operational parameters.

Dichloramine Formation

As the chlorine-to-ammonia weight ratio increases above 5:1 (up to ~7.6:1), or under lower pH conditions (pH 4.5 to 6.0), monochloramine reacts with additional hypochlorous acid to form dichloramine:

$\text{NH}_2\text{Cl} + \text{HOCl} \longrightarrow \text{NHCl}_2 \text{ (Dichloramine)} + \text{H}_2\text{O}$

Dichloramine produces strong, pungent, "swimming pool" odors and causes eye irritation. Operators actively avoid dichloramine formation in treated drinking water.

Trichloramine (Nitrogen Trichloride) Formation

At very high chlorine-to-ammonia weight ratios (> 7.6:1) or highly acidic conditions (pH < 4.4), dichloramine reacts further to produce trichloramine:

$\text{NHCl}_2 + \text{HOCl} \longrightarrow \text{NCl}_3 \text{ (Trichloramine)} + \text{H}_2\text{O}$

Trichloramine is an exceptionally noxious, volatile gas that causes severe lacrimation (tearing), respiratory irritation, and taste complaints even at concentrations below 0.05 mg/L.

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4. The Breakpoint Chlorination Curve

Breakpoint chlorination is the operational process of adding sufficient chlorine to water to completely satisfy initial inorganic demand, oxidize all ammonia and chloramine complexes, and achieve an uncombined, free available chlorine residual.

Chlorine Residual (mg/L)
   ^
   |                                                 / (Zone 4: Free Residual)
   |                                                /  (1:1 slope)
   |                 Peak (Zone 2)                 /
   |                   /\                         /
   |                  /  \                       /
   |                 /    \ (Zone 3)            /
   |                /      \                   /
   |  (Zone 1)     /        \                 /
   |  __________  /          \_______________/
   | /          \/            Breakpoint (Trough)
   +--------------------------------------------------------> Chlorine Dose (mg/L)
     Reducing     Combined     Chloramine      Free Available
     Demand       Formation    Oxidation       Residual

The Four Distinct Breakpoint Zones

1. Zone 1: Inorganic Demand / Immediate Destruction Chlorine reacts immediately with readily oxidizable inorganic reducing compounds, including ferrous iron ($\text{Fe}^{2+}$), manganous manganese ($\text{Mn}^{2+}$), hydrogen sulfide ($\text{H}_2\text{S}$), and nitrite ($\text{NO}_2^-$). These reactions consume chlorine completely; measured chlorine residual is zero.

2. Zone 2: Combined Residual Formation (Formation of Chloramines) Chlorine reacts with dissolved ammonia and organic nitrogen compounds. As the chlorine dose increases up to a weight ratio of approximately 5:1 $\text{Cl}_2:\text{NH}_3\text{-N}$, monochloramine and organochloramines form. The measured total chlorine residual rises steadily to a local maximum (the Peak). All measured residual is combined chlorine.

3. Zone 3: Oxidation and Destruction of Chloramines As additional chlorine is added beyond the 5:1 ratio, hypochlorous acid oxidizes the chloramines into nitrogen gas ($\text{N}_2$), nitrous oxide ($\text{N}_2\text{O}$), and hydrochloric acid ($\text{HCl}$): $2\text{NH}_2\text{Cl} + \text{HOCl} \longrightarrow \text{N}_2\uparrow + 3\text{HCl} + \text{H}_2\text{O}$ This destruction consumes both the added chlorine and the previously formed chloramines, causing the measured total residual curve to drop sharply until reaching the Breakpoint Trough (typically at a weight ratio of 7.6:1 to 10:1 $\text{Cl}_2:\text{NH}_3\text{-N}$). This zone is notorious for taste and odor complaints due to transient dichloramine and trichloramine intermediates.

4. Zone 4: Free Available Chlorine Residual Beyond the breakpoint trough, all ammonia and reducing agents have been fully oxidized. Any additional chlorine added exists purely as free available chlorine ($\text{HOCl}$ and $\text{OCl}^-$). The residual line rises in a direct 1:1 linear slope with applied dosage.

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5. Dosage, Demand, and Residual Calculations

The relationship between chlorine application and measurement is defined by the core mass-balance equation:

$\text{Chlorine Dose (mg/L)} = \text{Chlorine Demand (mg/L)} + \text{Chlorine Residual (mg/L)}$

$\text{Total Chlorine Residual} = \text{Free Available Residual} + \text{Combined Available Residual}$

Analytical Differentiation: The DPD Colorimetric Method

Operators utilize the N,N-diethyl-p-phenylenediamine (DPD) colorimetric method to quantify chlorine fractions:

• DPD-1 Reagent: Reacts instantly with Free Available Chlorine ($\text{HOCl} + \text{OCl}^-$) to produce a red/magenta color.
• DPD-3 Reagent (or Potassium Iodide): Added after DPD-1 to catalyze the reaction of chloramines, yielding the Total Chlorine Residual.
• Combined Chlorine: Calculated mathematically as $\text{Combined Residual} = \text{Total Residual} - \text{Free Residual}$.

Worked Example 3.1.1: Determining Breakpoint Chlorine Feed Rate

A water treatment plant treats 3.50 MGD with an ammonia nitrogen concentration of 0.25 mg/L $\text{NH}_3\text{-N}$ and an inorganic reducing demand of 0.80 mg/L. The target finished water free chlorine residual entering the clearwell is 1.60 mg/L. To ensure complete breakpoint chlorination, an operational ratio of 10:1 $\text{Cl}_2:\text{NH}_3\text{-N}$ is required.

Step 1: Calculate the chlorine required for ammonia oxidation: $\text{Chlorine for Ammonia} = 0.25\text{ mg/L } \text{NH}_3\text{-N} \times 10 = 2.50\text{ mg/L } \text{Cl}_2$

Step 2: Calculate total chlorine dosage required: $\text{Total Dose} = \text{Inorganic Demand} + \text{Ammonia Demand} + \text{Desired Free Residual}$ $\text{Total Dose} = 0.80\text{ mg/L} + 2.50\text{ mg/L} + 1.60\text{ mg/L} = 4.90\text{ mg/L}$

Step 3: Calculate daily chemical feed rate using the standard Pounds Formula: $\text{Feed Rate (lb/day)} = \text{Flow (MGD)} \times \text{Dose (mg/L)} \times 8.34\text{ lb/gal}$ $\text{Feed Rate} = 3.50\text{ MGD} \times 4.90\text{ mg/L} \times 8.34\text{ lb/gal} = 143.04\text{ lb/day of 100\% pure } \text{Cl}_2$