Identifying Specific Water Impurities
Identifying Specific Water Impurities
Many individual impurities can be quantified through water analysis techniques. Below is a discussion of most ionic individual contaminants.
The presence of calcium (Ca2+) and magnesium (Mg2+) ions in a water supply is commonly known as “hardness.” It is usually expressed as grains per gallon (gpg). Hardness minerals exist to some degree in virtually every water supply. The following table classifies the degree of hardness:
The main problem associated with hardness is scale formation. Even levels as low as 5 to 8 mg/L (0.3 to 0.5 gpg) are too extreme for many uses. The source of hardness is calcium- and magnesiumbearing minerals dissolved in groundwater. “Carbonate” and “noncarbonate” hardness are terms used to describe the source of calcium and magnesium. “Carbonate” hardness usually results from dolomitic limestone (calcium and magnesium carbonate) while “noncarbonate” hardness generally comes from chloride and sulfate salts.
Iron, which makes up 5% of the earth’s crust, is a common water contaminant. It can be difficult to remove because it may change valence states – that is, change from the water-soluble ferrous state (Fe2+) to the insoluble ferric state (Fe3+). When oxygen or an oxidizing agent is introduced, ferrous iron becomes ferric which is insoluble and so precipitates, leading to a rusty (red-brown) appearance in water. This change can occur when deep well water is pumped into a distribution system where it adsorbs oxygen.
Ferric iron can create havoc with valves, piping, water treatment equipment, and water-using devices. Certain bacteria can further complicate iron problems. Organisms such as Crenothrix, Sphaerotilus and Gallionella use iron as an energy source. These iron-reducing bacteria eventually form a rusty, gelatinous sludge that can plug a water pipe. When diagnosing an iron problem, it is very important to determine whether or not such bacteria are present.
Although manganese behaves like iron, much lower concentrations can cause water system problems. However, manganese does not occur as frequently as iron. Manganese forms a dark, almost black, precipitate.
Sulfate (SO42-) is very common. When present at lower levels, sulfate salts create problems only for critical manufacturing processes. At higher levels, they are associated with a bitter taste and laxative effect. Many divalent metal-sulfate salts are virtually insoluble and precipitate at low concentrations.
Chloride (Cl–) salts are common water contaminants. The critical level of chloride depends on the intended use of the water. At high levels, chloride causes a salty or brackish taste and can interfere with certain water treatment methods. Chlorides also corrode the metals of water supply systems, including some stainless steels.
Alkalinity is a generic term used to describe carbonates (CO32-), bicarbonates (HCO3–) and hydroxides (OH–). When present with hardness or certain heavy metals, alkalinity contributes to scaling. The presence of alkalinity may also raise the pH.
Although nitrate (NO3–) and nitrite (NO2–) salts may occur naturally, their presence in a water supply usually indicates man-made pollution. The most common sources of nitrate/nitrite contamination are animal wastes, primary or secondary sewage, industrial chemicals, and fertilizers. Even low nitrate levels are toxic to humans, especially infants, and contribute to the loss of young livestock on farms with nitrate-contaminated water supplies.
Chlorine, because of its bactericidal qualities, is important in the treatment of most municipal water supplies. It is usually monitored as free chlorine (Cl2) in concentrations of 0.1 to 2.0 ppm. In solution, chlorine gas dissolves and reacts with water to form the hypochlorite anion (ClO–) and hypochlorous acid (HClO). The relative concentration of each ion is dependent upon pH. At a neutral pH of 7, essentially all chlorine exists as the hypochlorite anion which is the stronger oxidizing form. Below a pH of 7, hypochlorous acid is dominant, and has better disinfectant properties than the anion counterpart. Although chlorine’s microbial action is generally required, chlorine and the compounds it forms may cause a disagreeable taste and odor. Chlorine also forms small amounts of trihalogenated methane compounds (THM’s), which are a recognized health hazard concern as carcinogenic materials. The organic materials with which the chlorine reacts are known as THM precursors.
In some cases, chlorine is also present as chloramine (i.e., monochloramine, NH2Cl) as a result of free chlorine reacting with ammonia compounds. The ammonia is added to a water supply to “stabilize” the free chlorine. Chloramines are not as effective a microbial deterrent as chlorine, but provide longer-lasting residuals.
This material is often produced on-site primarily by large municipalities via the reaction between chlorine or sodium hypochlorite and sodium chlorite. A more costly source of chlorine dioxide is available as a stabilized sodium chlorite solution. Chlorine dioxide has been used for taste and odor control and as an efficient biocide. Chlorine dioxide can maintain a residual for extended periods of time in a distribution system and does not form trihalomethanes (THM’s) or chloramines if the stabilized sodium chlorite form is used. The possible toxicity of the chlorate and chlorite ions (reaction byproducts) may be a concern for potable water applications.
Every water supply contains at least some silica (SiO2). Silica occurs naturally at levels ranging from a few ppm to more than 200 ppm. It is one of the most prevalent elements in the world. Among the problems created by silica are scaling or “glassing” in boilers, stills, and cooling water systems, or deposits on turbine blades. Silica scale is difficult to remove. Silica chemistry is complex. An unusual characteristic of silica is its solubility. Unlike many scaling salts, silica is more soluble at higher pH ranges. Silica is usually encountered in two forms: ionic and colloidal (reactive and nonreactive based on the typical analytical techniques). Silica can be present in natural waters in a combination of three forms: reactive (ionic), nonreactive (colloidal) and particulate.
Ionic Silica (reactive)
Ionic or reactive silica exists in an SiO2 complex. It is not a strongly-charged ion and therefore is not easily removed by ion exchange. However, when concentrated to levels above 100 ppm, ionic silica may polymerize to form a colloid.
Colloidal Silica (nonreactive)
At concentrations over 100 ppm, silica may form colloids of 20,000 molecular weight and larger, still too small to be effectively removed by a particle filter. Colloidal silica is easily removed with ultrafiltration, or can be reduced by chemical treatment (lime softening).
Colloidal silica can lower the efficiency of filtration systems (such as reverse osmosis). Any silica can affect yields in semiconductor manufacturing and is a major concern in high-pressure boiler systems.
Aluminum (Al3+) may be present as a result of the addition of aluminum sulfate [Al2(SO4)3] known as alum, a commonly used flocculant. Aluminum can cause scaling in cooling and boiler systems, is a problem for dialysis patients, and may have some effects on general human health. Aluminum is least soluble at the neutral pH common to many natural water sources.
The sodium ion (Na+) is introduced naturally due to the dissolution of salts such as sodium chloride (NaCl), sodium carbonate (Na2CO3), sodium nitrate (NaNO3) and sodium sulfate (Na2SO4). It is also added during water softening or discharge from industrial brine processes. By itself the sodium ion is rarely a problem, but when its salts are the source of chlorides (Cl–) or hydroxides (OH–), it can cause corrosion of boilers, and at high concentrations (such as seawater) it will corrode stainless steels.
Potassium is an essential element most often found with chloride (KCl) and has similar effects but is less common than sodium chloride. It is used in some industrial processes. The presence of KCl is typically a problem when only ultrapure water quality is required.
Most phosphates (PO43-) commonly enter surface water supplies through runoff of fertilizers and detergents in which “phosphates” are common ingredients. Phosphates also enter the hydrologic cycle through the breakdown of organic debris.
Phosphates are used in many antiscalant formulations. At the levels found in most water supplies phosphates do not cause a problem unless ultrapure water is required. Phosphates may foster algae blooms in surface waters or open storage tanks.