Sodium hypochlorite (\(\text{NaOCl}\)) is the active ingredient in liquid bleach, commonly used for sanitation and disinfection purposes. When this substance is introduced into water, it raises the water’s \(\text{pH}\) level, moving it toward the alkaline side of the scale. Commercial-grade sodium hypochlorite solutions typically have a high \(\text{pH}\) ranging between 11 and 13. This strong alkalinity is a direct result of the chemical composition. The initial rise in \(\text{pH}\) is a fundamental consideration for anyone using sodium hypochlorite for water treatment, such as in swimming pools or industrial cleaning applications.
The Chemistry Behind the pH Shift
The elevation in \(\text{pH}\) occurs because of a reaction called hydrolysis. In solution, sodium hypochlorite dissociates into sodium ions (\(\text{Na}^+\)) and hypochlorite ions (\(\text{OCl}^-\)). The hypochlorite ion acts as a weak base, attracting a hydrogen ion (\(\text{H}^+\)) from a water molecule (\(\text{H}_2\text{O}\)) to form hypochlorous acid (\(\text{HOCl}\)) and a hydroxide ion (\(\text{OH}^-\)).
The formation of the hydroxide ion directly increases the alkalinity of the water. \(\text{pH}\) is a measure of the concentration of hydrogen ions, and increasing the concentration of hydroxide ions decreases the hydrogen ion concentration, leading to a rise in \(\text{pH}\). Liquid sodium hypochlorite often contains excess sodium hydroxide (\(\text{NaOH}\)) from its production, which is added to stabilize the product. This residual sodium hydroxide also contributes to the initial high alkalinity.
Practical Implications of Elevated pH
High \(\text{pH}\) levels reduce chlorine effectiveness because the equilibrium shifts away from the desired compound. At high alkalinity, hypochlorous acid (\(\text{HOCl}\)), the powerful, fast-acting disinfectant, converts into the less effective hypochlorite ion (\(\text{OCl}^-\)). For example, \(\text{HOCl}\) makes up about 75% of total available chlorine at a \(\text{pH}\) of 7.0, but only about 25% at a \(\text{pH}\) of 8.0. This shift means a significantly larger amount of chlorine is required to achieve the same level of disinfection, leading to inefficient use of the product.
The high alkalinity also promotes the precipitation of mineral salts, primarily calcium carbonate, which manifests as scale. Scale can accumulate on surfaces, pipes, and equipment, potentially clogging filters and damaging sensitive components over time.
A highly alkaline environment can cause discomfort for users in applications like swimming pools. Water with an elevated \(\text{pH}\) (typically above 7.4 to 7.8) is known to cause irritation to the eyes and skin because it disrupts the natural \(\text{pH}\) balance of the body’s mucous membranes. High \(\text{pH}\) also contributes to cloudy water conditions, which affects the aesthetic and safety of the water system.
Strategies for pH Management
To counteract the \(\text{pH}\)-raising effect of sodium hypochlorite, acidic compounds must be added to the water to neutralize the excess hydroxide ions. This process involves introducing a controlled amount of acid to bring the overall \(\text{pH}\) back into an acceptable range, typically between 7.4 and 7.6 for sanitation purposes. The two most common chemicals used for this purpose are muriatic acid and sodium bisulfate.
Muriatic Acid
Muriatic acid, which is a diluted form of hydrochloric acid (\(\text{HCl}\)), is a potent liquid acid that effectively lowers both \(\text{pH}\) and total alkalinity.
Sodium Bisulfate
Sodium bisulfate (\(\text{NaHSO}_4\)) is another effective \(\text{pH}\) reducer, often sold as a dry, granular acid. This dry form is generally considered safer to handle than the highly corrosive liquid muriatic acid.
It is important to add the acid slowly and to thoroughly circulate the water before retesting, as adding too much acid can cause the \(\text{pH}\) to drop too low, which introduces a different set of problems.