Well and Pump Rehabilitation

Published On: May 30, 2017By Categories: Engineering Your Business

Part 1: Water chemistry basics

By Ed Butts, PE, CPI

While affecting a change to inferior water quality using an assortment of chemical agents once the water has reached the surface from a pumping well is quite feasible and common, the same usually cannot be said for correcting the natural state of the water in the well or aquifer.

In this case, it is necessary to understand the specific conditions of the water and well components in order to track possible changes over time to the production and efficiency of the well or pump. Over the next six columns, we will explore the various possible causes of declining well and pump efficiency and some of the methods used to identify and correct or abate these events.

I wrote a series of articles in 2004 and 2005 on water chemistry, corrosion, encrustation processes, and well rehabilitation procedures. This column—the first of three on the possible causes leading to the need for well or pump rehabilitation—will expand on all those discussions. We’ll go a little further into the nuts and bolts of corrosion and encrustation problems, the major problems leading to the need for well rehabilitation, and how to identify each of these conditions.

Part 2 will examine the specific physical, chemical, electrochemical, and biological processes involved in either plugging or corrosion (deterioration) of water wells and pumps and the water quality tests and indices used to identify potential or current problems in both.

Part 3 will wrap up this discussion with an overview of the various testing procedures used to gauge a well and pump’s performance and possible decline in each and some of the common treatment methods available.

The final three columns will detail some of the ways to determine and quantify well and pump efficiency and the common procedures used to obtain or maintain high operating efficiencies in new or used plants.

Once again, a few words of caution before we start. Although much of what I will write about is universal in concept and theory, you must apply this information to your specific area, groundwater quality, and well conditions. As we all know, one size never fits all. That is never as profound as when discussing problems with water wells.

Introduction to Chemistry

Figure 1. Each molecule of water consists of two atoms of hydrogen combined with one of oxygen.

Water is a basic compound in nature, each molecule consisting of two atoms of hydrogen combined with one atom of oxygen (Figure 1). As the “universal solvent,” water is capable of dissolving various types and concentrations of ions from natural substances it contacts.

Depending on the specific substance, concentration, and time of contact, the dissolution of this substance followed by absorption into the water molecule can have a negative, positive, or neutral impact on the relative quality of the water. This trait of the water molecule can and does occur in any type or state of water that exists: precipitation, surface water, or groundwater; solid, liquid, or vapor.

The primary factors impacting groundwater quality include the:

  • Chemical composition of the aquifer and dissolution from recharge water
  • Natural lithological and hydrogeological properties of the geologic unit
  • Various chemical processes and exchanges occurring within the geologic unit itself
  • Amount of time the water has remained in contact with the geologic unit and overlying formations (residence time).

All these factors and more can affect the type and quantities of dissolved constituents in a specific groundwater. The most abundant dissolved constituents measured in the Earth’s crust are the major ions, which can be both positively charged (cations) and negatively charged (anions). Because of the requirements of electroneutrality, cations and anions are present at equal concentrations in water and comprise most of the dissolved solids in groundwater.

The most abundant cations present in water are calcium (Ca), magnesium (Mg), sodium (Na), iron (Fe), and potassium (K). The most abundant anions are bicarbonate (HCO3), chloride (Cl), and sulfate (SO4).

By measuring the concentrations of these ions in groundwater samples, the ionic composition of the water is determined and the chemical quality of the water can be characterized and described. The ionic composition of water is used to classify it into ionic types and compounds based on the dominant dissolved cation and anion, expressed in milliequivalents per liter (meq/L). A milliequivalent (meq) is a measurement of the molar concentration of the ion, normalized by the ionic charge of the ion.

The dominant dissolved ion must be greater than 50% of the total. For example, water classified as a sodium bicarbonate type of water contains equal to or greater than 50% of the total cation milliequivalents as sodium and equal to or greater than 50% of the total anion milliequivalents as bicarbonate. If no cation or anion is dominant (greater than 50%), the water is classified as a “mixture” and the two most common cations or anions in decreasing order of abundance are used to describe the water type.

An example of this is a water containing cations of 40% sodium, 30% calcium, and 20% magnesium with anions of 52% bicarbonate, 30% sulfate, and 18% chloride would be classified as a sodium-calcium-bicarbonate-type of water.

When combinations of these anions and cations are present in concentrations exceeding the solubility product of the various species which may be formed, precipitates form as a result until the respective solubility products are no longer exceeded. This is when the water solution is balanced, forms a scale, or becomes corrosive. Consider when the concentrations of the calcium and sulfate ions exceed the solubility product of calcium sulfate, a solid phase (scale) of calcium sulfate forms as a precipitate.

Solubility products are exceeded for various reasons such as evaporation of the water phase; a change in pH, pressure, or temperature; and the introduction of additional ions which can form insoluble compounds with the ions already present in the solution. Generally, the most common cations impacting well water quality are iron, manganese, and calcium, which can readily combine with anions such as sulfate and carbonate to form scale-producing compounds including calcium carbonate (CaCO3), magnesium carbonate (MgCO3), and calcium sulfate (CaSO4).

As these reaction products precipitate on the surfaces of the well casing or screen, they form adherent deposits more commonly referred to as scale. As the scaling deposits increase, they eventually interfere with fluid flow, facilitate corrosive processes, or harbor bacteria.

Scale is an expensive problem for many water well systems, causing loss of well efficiency and yield and higher operating costs, eventually resulting in a well shutdown for cleaning and removal of the accumulated scale deposits.

To assist with determining water quality parameters, levels of specific cation and anions are often converted to equivalent calcium carbonate value. These, plus the formula weights, are shown in Table 1. There are many water quality parameters that can result in scale formation (encrustation) or corrosive conditions. To begin, let’s review the primary water chemistry factors involved with corrosion and encrustation: pH, alkalinity, iron and manganese, calcium compounds, and biofouling.


Arguably the most important single chemical characteristic of water is pH. A water’s pH is fundamental to virtually all chemical processes involving this compound. To review, there are three main groups of inorganic compounds: acids, bases, and salts.

An acid is any substance releasing hydrogen ions when it is mixed with water and is classified as either a strong or weak acid based on the amount of hydrogen ions released in water. Strong acids, such as sulfuric and hydrochloric acid, release a large concentration of hydrogen ions while weak acids, such as hydrogen sulfide, release few hydrogen ions.

Conversely, bases are substances producing hydroxyl ions when they dissociate (come apart) in water. Examples of bases include lime, caustic soda (sodium hydroxide), and soda ash (sodium carbonate). All these are chemicals commonly used in water treatment. As with acids, they can be subdivided into weak and strong bases. A strong base dissociates from water readily, resulting in the release of a high concentration of hydroxyl ions; weak bases, such as lime, dissociate poorly.

Salts are simply compounds resulting from an acid-base mixture. The process that results due to the mixing of an acid and a base to form a salt is known as neutralization.

The measurement of the concentration of acids and bases in water is performed using the pH scale (Figure 2). The pH scale runs from a low of 0 (most acidic) to a high of 14 (most basic). The scale is a logarithmic scale, which means every increment of pH measurement is ten times greater than the previous one. For example, a pH value of 6 is ten times more acidic than a pH value of 7 and so on in both directions.

Figure 2. The pH scale runs from most acidic (low of 0) to most basic (high of 14).

A pH value at the midpoint of the scale (7) is referred to as a neutral reading, comprised of a majority of neither an acid nor a base. For most waters used in a potable water system, the desirable pH range is between a low of 6.5 to a high of 8.5, although many other factors can affect the best pH for any given water.

Because a pH of 7 is called neutral, many people mistakenly believe it is the best value to maintain potable water. While in some cases this may be true, in many cases it is not. A pH value of 7 or even 7.5 may not be high enough to protect a piping system from corrosion if other protective elements in the water such as calcium hardness, iron, and manganese are not present. In this instance, water at a pH of 7 can cause an attack on the metallic components in a water system, ultimately resulting in a process called corrosion. In these cases, a pH level of 7.5 to 8.0 may be required to present a higher protective value.

Without a doubt, it all starts with pH. This is the reason many lead and copper programs or corrosion prevention systems use pH adjustment as the primary method of corrosion control. Conversely, if the pH of water (especially well water) is too high, a process known as precipitation can occur, causing iron, manganese, and calcium to come out of solution and deposit on the interior of pipe walls, resulting in scaling. In some extreme cases, the scaling of an internal pipe wall can be severe enough to completely close off the inside of a pipe or well screen slots in just a few years. In most cases, this can occur if the pH of the water is above 7.5 for iron and 8.3 for manganese, as well as other needed conditions like temperature and sufficient levels of iron or manganese are present.

There are many other situations where intentional elevation of pH values is needed to affect the removal of iron or manganese. In fact, in many water treatment systems the pH of the water is elevated to these values to accelerate oxidation and precipitation of iron and manganese so they will come out of solution rapidly and be filtered.

One last consideration of pH is for disinfection. Water disinfection, especially using chlorine, requires a value known as contact time for the chlorine to be in contact with the water to adequately kill or inactivate the pathogens in the water. In this case, a lower pH value will be more effective for chlorination than a higher pH. For example, to deactivate E. coli bacteria with a chlorine level of 0.05 mg/L will require nearly three minutes of contact time, while at a pH of 8.5 will require up to 20 minutes to affect the same deactivation level.

Systems requiring consideration of contact time to maintain adequate disinfection plus iron or manganese or corrosion control can really cause problems for engineers. If the pH of the water is naturally or artificially raised to 7.5–8.5 to control corrosion, an impact on contact time will occur. The same is true in the reverse instance where the pH of the water is lowered to 6.5–7 to improve disinfection as corrosion will likely occur.

Therefore, the key is to find some level of compromise between all aspects of water quality considerations. From this discussion, it should now be obvious why I began this section by implying pH was the single most important water quality consideration.


Alkalinity is a water quality element often overlooked, but actually important in the discussion of water chemistry.

It is basically the ability of a water to neutralize (or buffer) an acid. It occurs in three stages: hydroxyl, bicarbonate, and carbonate. The difference between the three stages is based on which one causes the alkalinity. The combined effect of all three types is reported as total alkalinity.

Alkalinity is important in water treatment as it has a huge impact on how much chemical is needed to raise or lower pH. For example, water with a high alkalinity (greater than 200 mg/L) will require a much greater volume of caustic soda to raise pH than it would at 30 mg/L This not only has a direct impact on the amount of chemical needed to change pH, but also affects the stability of pH.

In many cases with low alkalinity, it is difficult to maintain a consistent pH level in finished water, and specialized equipment is sometimes needed to monitor and adjust equipment used to change pH.

In conclusion, adequate knowledge of pH and alkalinity and the relationship between the two is critical for a complete understanding of corrosion control in a water system. Balancing the relationship of pH and alkalinity, along with using the appropriate chemicals, is crucial for optimum water system performance or well rehabilitation. Water with a pH level too high will usually cause scaling of pipe walls and well screen slots. Water with too low a pH level will often result in a corrosive attack onto metallic surfaces. The key is to work with all available information to strike a reasonable balance between the two.

Iron and Manganese

There is probably no water quality topic in our business discussed—and cussed—more than iron (Fe) and manganese (Mn). Although the two elements are both metals abundant in nature and many well waters, there are also many differences between them. While iron is everywhere and fairly easy to remove from water, manganese is a little rarer and usually much harder to remove.

Both elements are found in various states of chemical charge. Simple soluble iron is referred to as ferrous iron (Fe²+) and simple soluble manganese is called manganous manganese (Mn²+). When oxidized, iron changes to ferric iron (Fe³+) and manganese usually changes to insoluble forms: Mn³+ or Mn+.

Both elements are capable of forming combinations with many other elements and compounds. Although the presence of iron or manganese does not generally present any health problems, the aesthetic problems they cause due to staining of fixtures and taste and odor cause numerous complaints when they are present in their most common locations: groundwater supplies and surface water reservoirs.

Iron will generally cause complaints when the level is above 0.30 mg/L, while manganese causes problems at levels above 0.05 mg/L. Iron will generally cause reddish to brown stains; manganese will usually result in dark brown or even black stains in appearance. While both stains can be difficult to remove, the stains left from precipitated manganese can etch and damage fixtures and cause dark stains virtually impossible to remove.

Both elements can combine with microscopic living organisms to create a biofilm, known commonly as iron or manganese bacteria, which can be difficult to deal with and control.

In water treatment, iron and manganese are typically removed through one of two processes: ion exchange (water softening) or oxidation/reduction/filtration. Although water softeners are made to soften water through an exchange of sodium or potassium for calcium hardness, they can also work to remove low levels of iron and manganese with severe limitations.

Ion exchange uses a media (commonly called resin) the water percolates through and then a chemical exchange of sodium or potassium on the resin is exchanged for either iron or manganese.

Oxidation/reduction/filtration uses an oxidant such as air, chlorine, or ozone to change the chemical state of iron from a ferrous to a ferric state. In the ferric state, the iron has adequately reached the physical size and ability to be removed through a typical filter.

Iron left in the ferrous state or manganese in the manganous state is basically too attached to the water molecule and too small to be filtered and will travel through the system until encountering the atmosphere after discharge at a fixture. At this point, conditions of oxidation, temperature, and time are adequate to result in the change of state indicated above, resulting in precipitation and the stains often seen on fixtures.

In groundwater supply wells at adequate pH levels, the reaction between iron and manganese with the oxygen present in the wellbore can result in precipitation of the element, which can readily deposit iron oxide or manganese oxide onto the well’s surfaces and into well screen slots or perforations This reaction can be accelerated from the pressure drop of water incurred as water enters the wellbore.

Calcium Compounds

Next to the problems associated with iron and manganese compounds, the most common well scaling and plugging issue from minerals is generally the result of calcium carbonate (limestone) or calcium sulfate (gypsum) precipitation. Although precipitation of calcium compounds is much more prevalent in oil and gas wells due to the higher presence of these compounds and the higher temperature found in oil and gas, they are nonetheless also abundant in most water wells and can also result in significant scaling to the surfaces of well casing and drop pipe and openings of well screens and perforations.

Calcium carbonate is formed when elemental calcium ions (cations) react with carbonate ions (anions), while calcium sulfate forms from a reaction between calcium ions and sulfate ions.

The amount of scale that will precipitate out of water from each of these compounds is dependent on the temperature, pressure, and mineral content of the source water. As with iron and manganese, the principal determiners of the precipitation rate in a water well are tied to the temperature of the liquid, level of the specific contaminant, level of free and dissolved oxygen, time of exposure to oxygen, and the pressure drop of the water to convert from a soluble to a dissolved form.

One method of providing this pressure drop is through entrance velocity. As water enters the well through the well screen slots or perforations, there is often a drop in the associated vapor pressure of the fluid as it travels through the slot into the wellbore. This is a primary reason why many wells seem to endure an acceleration of a scaling or plugging action once the well screen begins to plug and the entrance velocity increases.

Because these variables often fluctuate in well water, predicting and controlling the type and amount of scale development can be difficult, requiring constant well monitoring and water quality analyses to account for these fluctuating parameters.

The most common parameter associated with calcium carbonate scaling potential is referred to as a water’s hardness. The calcium carbonate hardness of a water is measured in units of milligrams per liter (mg/L) or grains per gallon (gpg). Typically, a water’s hardness is considered acceptable if within a range of 60–150 mg/L (3.5–8.7 gpg). Water with a hardness level below 60 mg/L is generally corrosive to metal surfaces, while a level above 150 mg/L is usually excessively scale forming.


In addition to the potential well plugging problems caused from the precipitation of iron, manganese, and calcium compounds, another group of contaminants can seriously impact a well or pump’s performance and efficiency.

Biofouling, often simply referred to as iron bacteria, results from the microbial action occurring from microorganisms already present in the natural environment. These are live and dead organisms, their body parts and secretions, and other bacteria present in soil, the water, or on the surfaces of a well combined with metallic hydroxides—usually iron, manganese, or oxygen (air)—as an energy source.

The reaction between these constituents often leads to the development and deposition of a yellow, reddish, or brown gelatinous “slimy” organic film called a biofilm onto the surfaces, passages, and openings of wells and pumps.

Biofilms can be composed of many different types of organisms. The species present in a biofilm can vary in both location and time of year, heavily dependent upon the nutrients and amount of oxygen present in the water, the “seeding” or background microorganisms, the temperature and pH of the water, and the type of well casing, drop pipe, or well screen material.

The three principal nutrients needed for biofilm formation are assimilable organic carbon (AOC), nitrogen, and phosphorus. Introduction or presence of seed microorganisms is also necessary to initiate biofilm growth. Bacteria are nearly everywhere in the aqueous and natural environment and are usually present in vast numbers in raw or untreated well water.

Microorganisms from the genera Gallionella, Leptothrix, and Crenothrix are within the iron bacteria group and are the most common type found in well water. If surface water or soil enters a well, the bacteria may also be introduced and begin to thrive if conditions are favorable. If conditions in the well are suitable and the initial growth is not abated or controlled, the bacteria will quickly multiply in the structure of a biofilm.

Biofilm tends to be self-perpetuating as the initial growth of a biofilm results in increased development of more communities, leading to a generally progressively thicker layer of biofilm growth, also known as slime. The progression of growth from the initial development and attachment of a biofilm to surfaces begins within seconds of exposure and if left unabated, continues for months and sometimes years until the growth has reached a level where well or pump performance has been totally compromised.

In addition, non-pathogenic and pathogenic coliform organisms (including E. coli) and other various waterborne pathogens can enter the well from inadequately protected surface or shallow water sources and take up residence in the biofilm.

The presence of biofilm in a groundwater well can result in two separate serious issues. Biofilm can harbor and promote the growth of disease-causing bacteria, creating a threat to consumers of the water. Also, unchecked growth of a biofilm can drastically reduce the yield and efficiency of a water well or pumping plant in a short time, necessitating expensive and time-consuming downtime and rehabilitation procedures, which often are only marginally effective.

Biofilm in a water well can be difficult to detect and control as the microorganisms are so prevalent in nature and often adhere to the surfaces of drilling and pump installation tools and equipment, and can be transferred from one well to another if adequate precautions are not observed. Detection generally requires an ongoing examination of the potential for biofilm growth combined with monitoring a well’s yield as opposed to pumping water levels (specific capacity) to determine the well’s current biofilm condition. The present availability and access of downhole camera surveys can be used as an invaluable tool to assist in monitoring the biofouling of a well.


I hope I have conveyed the basic concepts associated with water quality parameters and how each applies to well corrosion or scaling. I have tried to give you enough basic water chemistry information to help provide guidance when you need to consider how each of these common conditions work in your specific application.

As I have said before, try and use the sciences when and wherever you can to help you in your day-to-day work. I think you’ll find they are not as scary as you may remember from high school—and can actually work for you.

Until next month, work safe and smart.

Ed Butts, PE, CPI, is the chief engineer at 4B Engineering & Consulting, Salem, Oregon. He has more than 40 years of experience in the water well business, specializing in engineering and business management. He can be reached at epbpe@juno.com.

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