Acid Addition Options for Commercial Pools
Control of pH in commercial swimming pools can require constant adjustment. Normal bicarbonate alkalinity provides buffering to help guard against wide fluctuations of pH in water. Increasing the bicarbonate alkalinity, however, increases the tendency for pH to drift upwards due to loss of dissolved carbon dioxide (CO2) to the air.[i]
With recommended levels of bicarbonate alkalinity in the water (>60 ppm) a normal pool pH of 7.2–7.8 cannot be maintained without a supersaturated concentration of dissolved CO2. Such high CO2 concentrations are not sustainable because CO2 is continuously lost from pool water to the atmosphere. This loss creates an upward drift in pH unless some form of acid is repeatedly added to offset the effect of CO2 loss. The Model Aquatic Health Code (section 4.7.3.2.8) addresses this tendency for pH drift by calling for the use of pH controllers in public swimming pools and spas.
The ACID-RITE® tablet erosion system, a new form of acid addition technology, makes the safety advantages of solid acid practical for commercial pools. This NSF 50 listed system allows for continuous acid addition, or as called for by a pH controller. Though traces of salt, common in sodium bisulfate, allow small amounts of hydrogen chloride gas to form, these fumes are minor when compared to those from muriatic acid. Details about this system can be found at www.acidrite.com.
The Acid-Rite system is a new choice in acid addition technology
Other acid types to consider to combat upward drift in pH
Carbon Dioxide: Since upward drift in pH results from loss of CO2 to the air, a direct approach to pH control is to re-inject carbon dioxide—from a compressed gas cylinder—to the pool water. This strategy conserves alkalinity. Yet, the downside of these acid additions to pool water is that they lower the pH by converting some of the bicarbonate in the water to carbon dioxide:
H+ + HCO3– → H2CO3 → H2O + CO2
This restores the [HCO3–]/[CO2] concentration ratio to a value corresponding to a lower pH, but in the process destroys bicarbonate. Eventually, alkalinity needs to be added to the water in one way or another. With CO2 addition instead of conventional acids, no bicarbonate is sacrificed. Alkalinity is conserved and future additions of bicarbonate are unnecessary.
This benefit comes with its own drawback. Addition of alkaline sanitizers (mainly hypochlorites), can result in a gradual build-up of alkalinity in the water. In addition, it is not uncommon for fill water to be highly alkaline. Any topping off with tap water amounts to a small alkalinity contribution. Water evaporates, but alkalinity does not. It simply accumulates until it is dealt with by acid addition. As alkalinity builds up, it becomes harder and less efficient to control pH by addition of carbon dioxide. Eventually, the elevated alkalinity will require application of a liquid or solid acid. Unless increases in alkalinity are addressed, carbon dioxide demand can climb steeply, making CO2 systems inefficient and costly.
Acidic disinfectants: Another option to lower pH is to simply use an acidic disinfectant, such as chlorine gas (pH 1) or trichlor (pH 3). This, however, will lead to an uncontrolled addition of acid and may well add too much, so that a pH increaser and/or alkalinity increaser may also need to be periodically added. Furthermore, there are other important considerations in the choice of disinfectant, such as the training and equipment required for safe handling of chlorine gas, or the risk of over-stabilization with chlorinated isocyanurates, like trichlor.
Apart from acidic sanitizers there are more mainstream forms of acid in both liquid and solid forms to consider.
Liquid (aqueous) acids such as muriatic acid and sulfuric acid are the most commonly used, due to price and availability. Here are some things to consider when using muriatic acid:
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Corrosion: Unlike sulfuric acid, muriatic acid releases corrosive fumes of hydrogen chloride gas. This gas is quite irritating as well as being corrosive to metallic equipment and fixtures in the pump room. Any place where chemicals are stored and used should be well ventilated. This is especially true where muriatic acid is involved.
Apart from the gradual damage that could happen due to muriatic acid fumes, any spray or spill of a concentrated acid could result in rapid damage to concrete or metal contacted by the leaked acid.
Safety Considerations: Because of the potential damage caused if a spray contacts eyes or skin, an eyewash fountain and safety shower are required by OSHA. The Model Aquatic Health Code (section 4.9.2.1.4) also calls for emergency eyewash stations in all chemical storage spaces. According to OSHA, eyewash stations must be located at a distance no greater than 10 seconds away from the location where strong acids (like muriatic acid or sulfuric acid) or oxidizers (such as chlorinating agents) are used.
Finally, risks associated with a strong acid spill require that secondary containment (such as diking or over-packs) be in place.
Solid acid, most commonly sodium bisulfate (aka sodium hydrogen sulfate), may present a safer, and in some cases more convenient, alternative to liquid acid. Solid acid emits little-to-no corrosive fumes, so pump-room corrosion is far less of an issue. Skin-contact hazards are also lower since there is no need to work with concentrated liquid acid. When acid addition is needed, the dry acid can be added to a large quantity of water, resulting in a much safer dilution. Since a solid spill is relatively easy to contain and clean up, no secondary containment is required.
Traditionally, dry-acid powders or granules are used for reducing pH and alkalinity in residential pools, by periodic manual addition. This is more cumbersome for commercial pools, where pH needs to be automatically controlled on a continuous basis. The recent introduction of Acid-Rite sodium-bisulfate tablet system, as mentioned above, has made application of sodium bisulfate easier and more convenient to apply.
[i] When bicarbonate is present the pH is given by the following equation:
where pK1 is a constant, approximately 6.4, depending on temperature;
[HCO3–] is the concentration of bicarbonate in the water; and
[CO2] is the concentration of dissolved carbon dioxide (and the trace concentration of carbonic acid) in the water.
