3.1 Chlorination Chemistry, Breakpoint Chlorination & Residuals

3.1 Chlorination Chemistry, Breakpoint Chlorination & Residuals

Disinfection is the single most critical barrier protecting public water systems from waterborne epidemics. Water treatment operators must possess an uncompromised mastery of aqueous chlorine reactions, the equilibrium kinetics of germicidal compounds, the breakpoint chlorination curve, and residual monitoring techniques.

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1. Commercial Forms of Chlorine in Water Treatment

Chlorine is delivered to municipal treatment plants in three primary chemical forms, each possessing distinct chemical strengths, storage imperatives, handling risks, and operational characteristics:

Operational Handling and Delivery Systems

• Chlorine Gas ($\text{Cl}_2$): Shipped in steel cylinders containing liquefied chlorine under ambient vapor pressure (approximately 70 to 100 psi at standard room temperature). The gas is 2.49 times heavier than air, meaning uncontained releases settle into pits, low-lying trenches, and basement pipe galleries. Evaporation of the liquid absorbs heat from the container walls; drawing gas too rapidly (more than 40 lb/day from a 150-lb cylinder or 400 lb/day from a 1-ton cylinder) causes the cylinder to freeze, reducing vapor pressure and interrupting feed. Ton containers feature two valves on the head: the top valve withdraws gaseous chlorine, while the bottom valve withdraws liquid chlorine for external evaporators.
• Sodium Hypochlorite ($\text{NaOCl}$): Often referred to as industrial bleach, commercial solutions are stabilized with sodium hydroxide (caustic soda) at a pH between 11 and 13. Over time—especially when exposed to temperatures above 85°F (29°C), direct sunlight, or transitional metals (iron, nickel, copper)—sodium hypochlorite naturally decomposes into sodium chlorate ($\text{NaClO}_3$) and oxygen gas. This gaseous off-gassing frequently causes positive displacement diaphragm metering pumps to become vapor-locked, necessitating degas valves or peristaltic pumps. Hypochlorite feed lines are also susceptible to calcium carbonate scaling at the injection quill where high-pH bleach contacts hard water.
• Calcium Hypochlorite ($\text{Ca(OCl)}_2$): Commonly marketed as High Test Hypochlorite (HTH), this solid compound is hygroscopic and must be isolated from all petroleum products, solvents, and moisture. For every 1.0 mg/L of chlorine applied as calcium hypochlorite, about 0.7 mg/L of calcium hardness (as $\text{CaCO}_3$) is added to the water. Operators prepare batch solutions in separate dissolving tanks before feeding the clarified supernatant liquid.

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2. Fundamental Aqueous Chlorination Chemistry

When gaseous chlorine is introduced into water via a vacuum-operated injector (ejector), it undergoes an instantaneous hydrolysis reaction:

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

This hydrolysis reaction is complete within tenths of a second at ordinary water temperatures. Each mole of gaseous chlorine yields one mole of hypochlorous acid ($\text{HOCl}$) and one mole of hydrochloric acid ($\text{HCl}$), consuming water alkalinity and driving finished water pH downward. Theoretically about 1.4 mg/L of natural alkalinity (expressed as $\text{CaCO}_3$) is consumed for every 1.0 mg/L of gaseous chlorine applied.

When sodium hypochlorite or calcium hypochlorite dissolves in water, hypochlorous acid is generated without producing hydrochloric acid:

$\text{NaOCl} + \text{H}_2\text{O} \rightleftharpoons \text{HOCl} + \text{Na}^+ + \text{OH}^-$

$\text{Ca(OCl)}_2 + 2\text{H}_2\text{O} \rightleftharpoons 2\text{HOCl} + \text{Ca}^{2+} + 2\text{OH}^-$

Both hypochlorite forms liberate hydroxide ions ($\text{OH}^-$), which modestly increases treated water pH.

Hypochlorous Acid Dissociation and the Decisive Role of pH

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

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

This equilibrium is entirely reversible and is governed by the equilibrium acid dissociation constant ($K_a \approx 2.9 \times 10^{-8}$ at 25°C, corresponding to a $pK_a$ of approximately 7.54). The concentration ratio between hypochlorous acid and hypochlorite ion depends almost exclusively on the pH of the water, with water temperature playing a secondary role:

Mechanism of Germicidal Action

Hypochlorous acid ($\text{HOCl}$) is an electrically uncharged, neutral molecule. Bacterial cell walls and viral capsids carry a net negative surface charge. Because $\text{HOCl}$ is uncharged, it easily penetrates the microbial cell membrane through passive diffusion. Once inside the pathogen, $\text{HOCl}$ oxidizes essential intracellular enzymes (particularly sulfhydryl groups in glucose-oxidizing enzymes), irreversibly damaging metabolic function and precipitating microbial cell death.

In contrast, the hypochlorite ion ($\text{OCl}^-$) carries a negative electrical charge. When an $\text{OCl}^-$ ion approaches the microbial surface, it encounters electrostatic repulsion from the negatively charged membrane, severely inhibiting its ability to penetrate the cellular barrier. Consequently, hypochlorous acid is approximately 80 to 100 times more potent as a disinfectant than the hypochlorite ion. At pH 7.0, an operator achieves disinfection in a fraction of the time required at pH 8.5 under identical disinfectant residual concentrations.

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3. The Chlorine Demand Curve & Breakpoint Chlorination

Chlorine is an aggressive, non-selective chemical oxidant. When introduced into raw surface or groundwater, it reacts with diverse organic and inorganic constituents before achieving an active free residual. The total chlorine required is dictated by the fundamental balance equation:

$\text{Chlorine Dose} = \text{Chlorine Demand} + \text{Chlorine Residual}$

• Chlorine Dose: The total concentration of chlorine added to the water stream, expressed in mg/L (or parts per million, ppm).
• Chlorine Demand: The concentration of chlorine consumed by reacting with reducing substances, organic matter, and ammonia in the raw water over a defined contact period, expressed in mg/L.
• Chlorine Residual: The total concentration of active chlorine compounds remaining in the treated water after the demand has been exerted, expressed in mg/L.

The Four Zones of Breakpoint Chlorination

Understanding the breakpoint chlorination curve is essential for optimizing primary disinfection, controlling objectionable chlorinous tastes and odors, and complying with water quality mandates:

Measured
Residual
  ^                                              / Zone 4: Free Residual
  |                                             /  (1:1 slope with dose)
  |                     Peak (Zone 2)          /
  |                       /\                  /
  |                      /  \                /
  |                     /    \              / 
  |                    /      \  Zone 3    /
  |                   /        \          / Breakpoint Dip
  |                  /          \________/
  |  Zone 1         /            (Minimum)
  |==============--/
  +-------------------------------------------------------> Chlorine Dose

1. Zone 1: Inorganic Demand Destruction (No Residual) When chlorine is initially dosed into raw water containing reduced inorganic substances—such as ferrous iron ($\text{Fe}^{2+}$), manganous manganese ($\text{Mn}^{2+}$), hydrogen sulfide ($\text{H}_2\text{S}$), and nitrite ($\text{NO}_2^-$)—these compounds immediately reduce hypochlorous acid to inactive chloride ions ($\text{Cl}^-$). In Zone 1, the measured chlorine residual remains exactly zero. No disinfection occurs.
2. Zone 2: Synthesis of Inorganic Chloramines (Combined Residual Rises) Once the immediate inorganic demand is satisfied, continued chlorine addition reacts with naturally occurring or agricultural ammonia nitrogen ($\text{NH}_3$). This produces combined chlorine compounds known as inorganic chloramines: $\text{NH}_3 + \text{HOCl} \rightleftharpoons \text{NH}_2\text{Cl}\text{ (Monochloramine)} + \text{H}_2\text{O}$ $\text{NH}_2\text{Cl} + \text{HOCl} \rightleftharpoons \text{NHCl}_2\text{ (Dichloramine)} + \text{H}_2\text{O}$ $\text{NHCl}_2 + \text{HOCl} \rightleftharpoons \text{NCl}_3\text{ (Trichloramine / Nitrogen Trichloride)} + \text{H}_2\text{O}$ In Zone 2, the measured total chlorine residual increases linearly in proportion to the applied dose, but this residual is entirely combined available chlorine, dominated by monochloramine at neutral to alkaline pH. While monochloramine is a recognized disinfectant, it acts much slower than free chlorine.
3. Zone 3: Chloramine Destruction and Oxidation (Residual Drops) As the chlorine-to-ammonia weight ratio surpasses roughly 5:1, additional hypochlorous acid begins to oxidize the chloramines formed in Zone 2. This complex sequence converts chloramines into inert nitrogen gas ($\text{N}_2$), nitrous oxide ($\text{N}_2\text{O}$), and hydrochloric acid: $2\text{NH}_2\text{Cl} + \text{HOCl} \longrightarrow \text{N}_2\uparrow + 3\text{HCl} + \text{H}_2\text{O}$ As chloramines are destroyed, the measured chlorine residual steadily declines, creating the descending slope of the curve. During Zone 3, water exhibits intense, unpleasant "swimming pool" odors and causes eye irritation due to volatile dichloramine and trichloramine off-gassing.
4. Zone 4: Breakpoint and Free Residual Development The lowest point on the curve is the Breakpoint. At this precise juncture, all ammonia nitrogen has been completely oxidized. The theoretical stoichiometric ratio required to reach breakpoint is 7.6:1 (parts chlorine to parts ammonia-nitrogen by weight); in field operations, raw water organic matter typically elevates this practical demand to between 8:1 and 10:1. Once the breakpoint is reached, any additional chlorine dosed is unhindered by demand reactions and yields an equivalent, 1:1 increase in Free Available Chlorine Residual ($\text{HOCl} + \text{OCl}^-$).

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4. Chlorine Residual Measurement & Quality Control

To ensure regulatory compliance and guarantee microbiological safety without creating severe aesthetic defects, operators must accurately distinguish between different residual species.

Analytical Fractions of Chlorine Residual

• Free Available Chlorine: The combined concentration of hypochlorous acid ($\text{HOCl}$) and hypochlorite ion ($\text{OCl}^-$). It is the fastest-acting, most effective germicide.
• Combined Available Chlorine: Chlorine combined chemically with ammonia (chloramines) or organic nitrogen compounds (chloro-organics). It is a weaker disinfectant with slower reaction kinetics.
• Total Chlorine: The mathematical sum of all active chlorine species in the water: $\text{Total Chlorine} = \text{Free Chlorine} + \text{Combined Chlorine}$.

The DPD Colorimetric Method (Standard Methods 4500-Cl G)

The primary field and laboratory test for drinking water utilities is the N,N-diethyl-p-phenylenediamine (DPD) colorimetric method:

1. Free Chlorine Procedure: A treated water sample is buffered to a pH between 6.2 and 6.5 and mixed with DPD indicator reagent. Free chlorine instantly oxidizes the DPD amine to form a magenta-red meriquinoid dye (Wurster dye). The absorbance is read photometrically or matched against calibrated color discs within 1 minute of reagent addition to prevent monochloramine breakthrough.
2. Total Chlorine Procedure: Potassium iodide ($\text{KI}$) is introduced into the sample along with DPD. Combined chlorine compounds (monochloramine and dichloramine) catalytically oxidize the iodide ions into elemental iodine ($\text{I}_2$). The liberated iodine then reacts quantitatively with DPD to generate the magenta dye. The sample is allowed to incubate for two minutes before photometric reading.
3. Combined Chlorine Determination: Combined chlorine cannot be measured directly on a standard colorimeter; it is determined by subtraction: $\text{Combined Chlorine (mg/L)} = \text{Total Chlorine (mg/L)} - \text{Free Chlorine (mg/L)}$

Common Testing Interferences & Troubleshooting

• Monochloramine Breakthrough: If an operator delays reading the free chlorine sample beyond 60 seconds, monochloramine begins to slowly react with DPD, artificially inflating the apparent free chlorine measurement.
• High Chlorine Bleaching: At free chlorine concentrations exceeding 10.0 mg/L, the DPD reagent is oxidized beyond the pink Wurster dye stage into a colorless imine compound. An operator testing highly chlorinated water (such as newly disinfected distribution mains) might observe an initial flash of red that instantly turns crystal clear, incorrectly reporting zero chlorine. When high concentrations are suspected, the sample must be diluted with chlorine-demand-free deionized water before testing.
• Oxidized Manganese ($\text{Mn}^{4+}$): Manganese dioxide oxidizes DPD directly, producing a false-positive free chlorine reading. Operators can correct for this interference by quenching chlorine with sodium arsenite or thioacetamide in a blank duplicate sample.

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5. Practical Dosing Calculations & Worked Examples

Water operators frequently calculate feed rates to maintain required dosages under varying plant flow conditions.

The Fundamental Pounds Formula

$\text{Chemical Feed Rate (lb/day)} = \text{Flow Rate (MGD)} \times \text{Dose (mg/L)} \times 8.34\text{ lb/gal}$

Where $8.34\text{ lb/gal}$ is the mass of one gallon of water, converting flow in millions of gallons per day (MGD) and concentration in milligrams per liter (parts per million) directly into pounds per day.

Worked Example 1: Gaseous Chlorine Feed Rate

A conventional water treatment facility processes an average finished flow of 3.8 MGD. Laboratory jar tests and historical records establish that the raw water exerts a chlorine demand of 2.3 mg/L. The state operating permit mandates a free available chlorine residual of 1.7 mg/L leaving the clearwell.

Step 1: Calculate the required chlorine dose in mg/L. $\text{Dose} = \text{Demand} + \text{Residual} = 2.3\text{ mg/L} + 1.7\text{ mg/L} = 4.0\text{ mg/L}$

Step 2: Calculate the required daily chlorine gas feed rate. $\text{Feed Rate (lb/day)} = 3.8\text{ MGD} \times 4.0\text{ mg/L} \times 8.34\text{ lb/gal} = 126.77\text{ lb/day}$

The chlorinator rotameter should be set to deliver approximately 126.8 lb/day of chlorine gas.

Worked Example 2: Liquid Sodium Hypochlorite Solution Feed Rate

A groundwater well operates at a flow rate of 850 gallons per minute (gpm). The required chlorine dosage is 2.2 mg/L. The utility feeds commercial sodium hypochlorite solution containing 12.5% available chlorine by trade percent (density = 10.0 lb/gal; active available chlorine = 1.25 lb $\text{Cl}_2$ per gallon of solution).

Step 1: Convert the flow rate from gpm to MGD. $\text{Flow (MGD)} = \frac{850\text{ gpm} \times 1440\text{ min/day}}{1,000,000} = 1.224\text{ MGD}$

Step 2: Calculate the required pounds of 100% chlorine per day. $\text{Pounds } \text{Cl}_2\text{/day} = 1.224\text{ MGD} \times 2.2\text{ mg/L} \times 8.34\text{ lb/gal} = 22.46\text{ lb/day}$

Step 3: Calculate the gallons of 12.5% hypochlorite solution required per day. $\text{Solution Feed Rate (gal/day)} = \frac{22.46\text{ lb/day } \text{Cl}_2}{1.25\text{ lb } \text{Cl}_2\text{/gal solution}} = 17.97\text{ gal/day}$

Step 4: Calculate the metering pump output in milliliters per minute (mL/min). $\text{Pump Rate (mL/min)} = \frac{17.97\text{ gal/day} \times 3785\text{ mL/gal}}{1440\text{ min/day}} = 47.23\text{ mL/min}$

The chemical feed diaphragm pump should be calibrated to deliver 47.2 mL/min.