Water Chemistry Fundamentals


Water is one of our most important resources. It is essential for all plant and animal life. Therefore, it is not so surprising that 70% of the earth is covered by water. The earth has 326 cubic miles of water. However, 97% of the water supply is contained in the salty oceans. Only 3% of the water is fresh. And of the total fresh water supply, 90% is locked up in the frozen ice caps and glaciers of Antarctica and Greenland. When viewed this way, our fresh water supplies are a valuable commodity.

Americans use about 70 gallons of fresh water per person per day in the home. Industry and commercial businesses use the most water, however. Manufacturing plants use about 140 billion gallons of water every day. Fortunately, only 2% of this water is consumed, the remaining 98% is returned to the water supply as industrial waste water. It takes about 270 tons of water to make a ton of steel, 250 tons of water to make a ton of paper, and about 10 gallons of water to refine 1 gallon of gasoline.

In this section on Fundamentals of Water Chemistry we will review the following topics

  • Water as a resource

  • Typical industrial water sources

  • Basic water analysis

  • Methods of reporting water analysis results

  • Water chemistry terms

  • Acids and bases

  • pH

  • Conductivity

  • Optical and electronic spectroscopy


Industry draws its water from surface, groundwater and municipal water supplies. Surface waters come from streams, ponds, lakes and reservoirs. Ground waters are pumped from wells, mines and springs. Ground water typically contains higher concentrations of dissolved substances than surface waters because ground water has percolated through rock and soil formations. As it flows through the ground, it dissolves mineral deposits into the water. Because of the filtering action of the ground, however, ground waters are lower in suspended matter than surface waters. By contrast, surface waters are higher in suspended solids, color and runoff pollutants.

The chemical characteristics of ground waters are influenced by the geology of the surrounding strata. These include limestone (calcium carbonate), sandstone (silica), and gypsum (magnesium sulfate). Ground waters may also contain iron, manganese, sulfate-reducing bacteria, and iron bacteria.

The upper layers of the earth’s crust contain residual vegetable and animal matter, along with bacteria and other microorganisms. As a result, surface waters contain some organic matter and microorganisms. The composition of surface water is subject to rapid change because of rainfall and pollution. The most troublesome impurity in surface water is colloidal matter. This is extremely small particles of normally insoluble substances such as silica, alumina and hydrated oxides of iron in suspension. Coagulation, sedimentation and filtration are required to remove these impurities from water.

Odor and taste are often contributed by sulfides, microorganisms or industrial contamination. Color has its source in natural organic contaminants such as lignins, humic acids and tannins.

Microorganisms include bacteria, algae growths, slimes, diatoms, protozoa and mollusks. These organisms are responsible for deposits and corrosion in water supplies and cooling water systems.

Municipalities draw their water from surface and groundwater supplies. The City of Chicago, for example, operates the world’s largest water purification plant on the shores of Lake Michigan. It treats over 1.75 billion gallons of water per day. Typical municipal treatment includes filtration for suspended solids removal and chlorination to kill harmful bacteria. Other treatments include zinc and phosphate for corrosion control and fluoridation to minimize dental cavities.


The average layman judges the quality of water based on its taste, color and odor. If the tap water tastes good and has no color or odor, it is assumed to be of “good” quality and fit for human consumption. Water used for industrial or commercial purposes must also be of acceptable quality. As we have seen, however, it is necessary to know more about water quality than just taste, color and smell. Will the water form scale in heat transfer equipment? Will it cause corrosion of system metal? Or does it contain microorganisms that can foul plant equipment? To answer these questions we need to understand basic water analysis and chemistry.

The following is a typical water analysis

Total Hardness as CaCO3

142 ppm

Calcium Hardness as CaCO3

90 ppm

Magnesium Hardness as CaCO3

52 ppm

P alkalinity

0 ppm

M alkalinity

110 ppm

Chloride as NaCl

28 ppm

Sulfate as SO4

5 ppm

Silica as SiO2

5 ppm




300 umhos/cm

Hardness relationships

The hardness is defined as the calcium and magnesium content of the water. We are concerned about water hardness because these impurities are the ones that cause scale to form in plant equipment. In this case, the calcium hardness is 90 ppm and the magnesium hardness is 52 ppm. The combined or total hardness is 142 ppm.

In practice the total hardness is measured by direct titration with standard EDTA solution. The calcium hardness is also determined by direct titration. The magnesium hardness is calculated as the difference between total hardness and calcium hardness.

Alkalinity relationships

The alkalinity of the water is defined as the total of the bicarbonate (HCO3), carbonate (CO3) and hydroxide (OH) concentrations. In actual practice, the alkalinity is determined by titration with N/50 sulfuric acid to the phenolphthalein or ‘P’ alkalinity endpoint at pH 8.3 and to the methyl purple (or methyl orange) indicator endpoint at pH 4.3. The P and M alkalinities are indicated in the water analysis report. This information is used to calculate the specific alkalinity based on the relationships indicated in the following chart.

Alkalinity Relationships

























In our example, the P alkalinity is zero, so all alkalinity is present as bicarbonate alkalinity which equals M. In practice, only two of the three forms of alkalinity can exist at the same time. It is possible to have bicarbonate alone, bicarbonate and carbonate, carbonate alone (a rare occurrence), carbonate and hydroxide, or hydroxide alone. Stated differently, P alkalinity is the measure of all hydroxides and ½ the carbonates. M alkalinity measures all hydroxides, carbonates and bicarbonates; hence M alkalinity is referred to as total alkalinity.


Units of Measurement

Results of water analyses are expressed in several ways. The most precise are those that state the weight per unit volume, such as milligrams per liter (mg/l) or grains per gallon. Less precise are the measuring units of parts per million (ppm) or parts per billion (ppb), which can be related to either weights or volumes. In water analysis, one liter weighs 1000 grams, so mg/l and ppm are used interchangeably.

The older designation of grains per gallon is still in use by some municipalities, and it is commonly used in ion exchange calculations. 1 grain per gallon is equal to 17.12 ppm.

Useful Conversion Factors

Grains per gallon x 17.12 = ppm (or mg/l)

Grains per gallon x 142.86 = pounds per million pounds

US gallon x 1.2 = Imperial gallon

1 gram = 15.43 grains

1 pound = 7000 grains = 453.6 grams

1 gallon = 3.785 liters

1 cubic foot of water = 7.48 US gallons

1 pound per 1000 gallons = 120 ppm


Substances dissolved in water form charged species called ions. The fundamental building blocks of ions are atoms. Each of these building blocks has its own atomic weight that remains unchanged whenever chemical combinations are made with other atoms. When more than one of these atoms is combined, the resulting substance is called a molecule.

Ionization, Cations and Anions

When salts and other inorganic compounds are dissolved in water they no longer exist in the form of compounds. Rather, they dissociate (break up) into positively or negatively charged ions that conduct current. The electrical conductivity of water is an indication of the amount of dissolved ions present, and the higher the conductivity, the higher the dissolved solids content.

The presence of ions, their ability to conduct current, and the effect of increasing concentrations of salts in solution can be demonstrated readily by using a current source in series with two electrodes and a light bulb. Almost no current can flow when the electrodes are immersed in distilled water. However, if a salt, such as sodium chloride, is added, the molecules dissociate into positively and negatively charged ions that do conduct current and the electric bulb will light. Increasing the salt content will cause the light to burn brighter, since the ability to conduct current is a function of total electrolyte content. Certain gases, such as ammonia and carbon dioxide, also produce current-carrying ions in solution.

If two metal strips placed in a solution of ionizable substances are connected to the positive and negative poles of a battery, the ions that go to the negative pole (cathode) are the positive ions. Positive ions are called cations. Negatively charged ions are called anions because they are attracted to the positive pole (anode). Magnesium, calcium, sodium, aluminum and iron ions are attracted to the negative cathode, and are therefore cations. Carbonate, sulfate, and chlorides are attracted to the positive anode, and are thus anions.

The atoms and molecules of compounds are bonded together by the attraction of the positive and negative charges on the ions. Some atoms have a single positive or negative charge, such as sodium which has one positive charge or chloride which has one negative charge. Calcium has two positive charges and carbonate has two negative charges. Atoms combine with one another so as to maintain the electrical neutrality of the compound. Thus calcium with two positive charges would combine or link up with one carbonate ion which has two negative charges, or alternatively two chloride ions which have one negative charge a piece. These chemical bonds or linkages are called valences. The valence is the atomic charge, either positive or negative, on the atom or molecule.

Each atom has a specific atomic weight. Atoms that combine together to form molecules result in substances that have a specific molecular weight. The equivalent weight of a substance refers to the relationship between the atomic weight or molecular weight and the valence. Equivalent weight is the molecular weight divided by the valence. For example, oxygen has two negative charges and an atomic weight of 16. The molecular weight or 16 divided by the valence of 2 give an equivalent weight of 8.

Calcium carbonate (CaCO3) has a molecular weight of 100. When calcium carbonate dissolves in water it forms calcium (Ca+2) and carbonate (CO3-2). The valence is 2, so the equivalent is 100 divided by 2 or 50.

Because of the convenience of working with a molecular weight of 100 and an equivalent weight of 50, calcium carbonate is widely used as a reference basis for equivalent weights in water chemistry. The concentration of any substance in terms of CaCO3 is obtained by multiplying the number of equivalents of the substance found by 50 (the equivalent weight of calcium carbonate). The number of equivalents of any substance is determined by dividing the concentration of the atom or molecule in mg/l (or ppm) by its equivalent weight. For example, magnesium (Mg+2) has an atomic weight of 24.3 and an equivalent weight of 12.15. If a water sample contains 48.6 mg/l magnesium, it has 48.6/12.15 = 4 equivalents of magnesium. The concentration of magnesium expressed in terms of calcium carbonate (CaCO3) is 4.0 equivalents times 50, the equivalent weight of calcium carbonate. The magnesium concentration is 200 mg/l expressed as calcium carbonate.

Equivalent weight is a useful basis for evaluating the accuracy of water analyses. In any properly balanced water analysis, the cations (positive charge) must equal the negatively charged anions when both are expressed in terms of equivalents. To find the number of equivalents, divide the concentration of the substance (mg/l) by the equivalent weight. Thus 80 mg/l of calcium expressed as calcium ion divided by the equivalent weight of calcium (20) yields 4 equivalents. The calcium concentration expressed as calcium carbonate is 4 equivalents times 50, the equivalent weight of calcium carbonate, or 200 mg/l as calcium carbonate. In this way, convert all cations and anions listed in the water analysis to equivalents. If the analysis is off by more than 10%, check for error or omissions in the analysis.


An acid is defined as any compound that dissociates to produce hydrogen ions (H+) in solution. A base is any compound that gives hydroxyl ions (OH) in solution. A strong acid yields more hydrogen ions in solution that a weaker acid. Likewise, a strong base produces more hydroxyl ions than a weak base.

When an acid reacts with a base, water and a salt are formed. For example, sulfuric acid (H2SO4) added to sodium hydroxide (NaOH) yields sodium sulfate (Na2SO4) and water (H2O). Sodium sulfate, when redissolved, forms two sodium ions (Na+) and one sulfate acid radical (SO4-2).

All inorganic acids are combinations of hydrogen ions and a neutral salt. Nitric acid (HNO3), hydrochloric acid (HCl) and sulfuric acid (H2SO4) are examples.


Water dissociates or ionizes to form [H+] ions and [OH] ions. The hydrogen ion plays an important role in many chemical reactions such as acid-base titration, precipitation reactions and oxidation-reduction equilibrium. In addition, it affects the corrosion rate on system metals, determines the potential for scale deposition and influences the environmental impact of wastewater discharges.

At equilibrium, the concentration of hydrogen ion times the concentration of the hydroxide ion is a constant value, known as the dissociation constant. For water, the dissociation constant, K, is 10-14. (10-14 is scientific notation for 0.000 000 000 000 01 or 1 with 14 decimal places to the left.) The dissociation equation is represented as

K = [H+] X [OH] = 10-14

If the hydrogen ion concentration is 10-7, the OH concentration must be 10-7 to maintain the dissociation constant, K. If the hydrogen ion concentration increases to say, 10-5, the OH concentration decreases to 10-9.

It is cumbersome to refer to hydrogen and hydroxyl ion concentrations as 10-7 or 0.000 000 1. To simplify matters, the hydrogen ion concentration is expressed as the negative logarithm. In this way, the negative logarithm of 10-7 becomes 7. The abbreviation for the negative logarithm of the hydrogen ion concentration is pH. Likewise, the negative logarithm of the hydroxyl ion concentration is pOH.

pH is expressed on a scale of 0 to 14. pH 7 is the midpoint of the scale and denotes neutrality. As the pH decreases, the hydrogen ion concentration increases and the solution becomes more acidic. As the pH increases, the hydrogen ion concentration decreases causing the solution to be less acidic or more basic.

As indicated by the logarithmic pH scale, a shift in pH of 1 unit represents a tenfold increase or decrease in hydrogen ion concentration. A two unit shift represents a hundredfold change. These changes are caused by the addition of acids and bases to the water. Theoretically, pure water, such as distilled or deionized water, has a pH of 7.0. Carbon dioxide gas, however, can cause the pH to be 6.5 or lower. Other impurities can have a more or less dramatic affect on the pH.

In natural water, the carbon dioxide/carbonate/bicarbonate alkalinity equilibrium determines and controls the pH of the water. In water chemistry, this equilibrium is measured and reported a P-alkalinity and M-alkalinity. The pH value of a solution defines only the degree of acidity or basicity as related to [H+] and [OH] ions, and makes no reference to the total quantity of acid or alkali in solution. Likewise, the determination of total alkalinity or acidity of a solution by titrimetric methods does not necessarily reflect the degree of acidity or basicity. Hence, there is no direct mathematical relationship between P and M alkalinity, and pH.

pH of Common Solutions

Acidic Solutions



Soft drinks


Lemon Juice


Basic Solutions

Sodium bicarbonate


Milk of magnesia


Household ammonia


Measuring pH

Color indicators like phenolphthalein and methyl purple show a distinct color change with changes in pH. Other indicators have similar properties, but change color at different pH ranges. pH indicator solutions can be used to determine pH in buffered water with a fair degree of accuracy. A small amount of indicator is added to the sample and the resultant color compared to a set of known standards. The closest color match is referenced as the pH of the test solution.

A pH meter is commonly used in the laboratory to measure pH. Simply, a pH meter is a millivolt meter that measures the electric potential difference between a pH glass electrode and a saturated calomel reference electrode pair. The millivolt (mV) signal varies linearly with pH according to the following equation.

mV = -59.16(pH – 7.000)

Under ideal conditions at 25 0C, the mV signal would be 0 at pH 7.000. Real-life electrodes generate imperfect signals, however, with mV = 0 within a range of pH 6.2 to 7.8, and the slope varying as much as 8% less than 59.16. For these reasons, the pH meter must be calibrated with appropriate pH standards to determine the precise slope and intercept for that particular meter and electrode pair.


Dissolved ions conduct electric current. Conductivity, or specific conductance, is the measure of the water’s capability for conducting this electric current. The more dissolved ions in solution, the higher the conductivity and visa versa.

The basic unit of conductivity measurement is the mho, now more commonly called the siemen. (The term siemen honors two brothers, Werner von Siemens and Sir William Siemens, who were 19th century electroplaters and industrialists.) One mho is equivalent to 1 siemen. Specific conductance is the reciprocal of electrical resistance, which is measured in ohms.

The mho is too large for most water analysis work. Therefore, the conductance is reported in microsiemens or millisiemens. 1 siemen = 1000 millisiemens = 1,000,000 microsiemens. Because absolutely pure water is devoid of dissolved ions, the conductivity is very low and the resistance is very high. High purity water has a conductivity of 0.055 microsiemens per cm. Compare this to typical Chicago drinking water, for example, having a conductivity of around 300 microsiemens/cm, or seawater at 53,000 microsiemens/cm.

It is cumbersome to report conductivity values of high purity water in these low ranges. For this reason, the purity of demineralized water is more commonly reported in units of resistivity. Resistivity is the inverse of conductivity, and is reported in megohms per centimeter. One megohm is equal to 1 million ohms.

Measuring Conductivity

Conductivity is measured by applying an electric potential across two conductivity electrodes (plates) immersed in a test solution. The conductivity is determined from the voltage and resultant current produced within this conductivity cell.

The conductivity measured by this method is affected by the cell geometry. The size of the electrodes and their distance apart affect the electric current. For this reason, specific conductance is used to define the conductivity of the test solution. This standardizes the conductivity measurement and compensates for any differences in cell geometry.

The specific conductivity is determined by multiplying the measured conductivity by the cell constant. The cell constant is calculated by dividing the length of the distance between the electrodes (L) by the area of the electrodes (A). This value is supplied by the probe manufacturer.

Specific conductivity = Measured conductivity x L/A

Conductivity is affected by the temperature of the sample. These temperature effects differ, depending on the type of solution being tested. Many conductivity meters have automatic temperature compensation circuitry that references the readings to a standard temperature of 25 oC.

Conductivity/TDS Relationships

Conductivity measurements are frequently used to estimate the total dissolved solids content of water. But the concentration and activity of the ions present in the sample will affect the specific conductivity of the solution. If the test solution is primarily sodium chloride (NaCl) the TDS will be 50% of the specific conductance. If the water sample consists of a blend of chlorides, bicarbonates, and sulfate, however, the TDS will be between 65% and 80% of the conductance. Different types of conductivity standards are available to calibrate conductivity meters. Standard solutions of sodium chloride and potassium chloride, or a mixture of 40% sodium sulfate, 40% sodium bicarbonate and 20% sodium chloride are commonly used. The latter is called a “442 solution.” Multiplying the specific conductance by 0.50 gives the TDS as NaCl solution. Multiplying by 0.65 gives the TDS as 442 solution. For most water analysis work, the 0.65 conversion factor seems appropriate. For unneutralized boiler water, containing highly conductive hydroxide alkalinity, the factor is approximately 0.5.


Certain water impurities will react with specific test reagents to produce a distinct color. Chlorine, for example, reacts with DPD reagent to produce a characteristic magenta color. The intensity of the color is in direct proportion to the chlorine concentration, i.e. the higher the chlorine concentration, the deeper the color. This same principle is used in specific tests to determine the concentration of many other ions including iron, silica, molybdenum, manganese, phosphate, zinc and many others.

The intensity of the sample color is measured by visually comparing the color of the test sample to a series of standards. Various test equipment is available to accomplish this task. They consist of color wheels and comparator blocks that make matching of the sample with the correct standard simple and accurate. These test kits include all the necessary reagents for conducting the specific test along with detailed instructions for obtaining accurate results.

Some individuals have difficulty matching the color of the sample with the standards. This is particularly true if you are color blind. If this is the case, or if a more precise and accurate measurement is required, the color intensity can be measured electronically with a photometer (sometimes called a photoelectric colorimeter, or just a colorimeter). This instrument can be used for absorption and emission measurements depending on the test method.

Other instruments for spectroscopy measurements include spectrophotometers. These offer a wide range of test methods, come preprogrammed from the manufacturer and include all necessary test reagents for obtaining fast, reliable and accurate measurements.

The criteria for selection of an optical colorimeter, photometer, or spectrophotometer depend on the substance being tested, the level of precision and accuracy required, and the skill level of the technician. In some cases, the EPA dictates the test method required for sampling and testing. Further guidance is available from the manufacturer or your consultant.


Determining the quality of water for industrial applications goes beyond color, odor and taste. Key impurities include hardness, alkalinity, specific cations and anions, pH, and conductivity.

Analyzing a water sample for its impurity level is the first step in determining its suitability for the intended purpose. Fortunately, this task is made easier by test equipment such as titration burettes, pH meters, conductivity meters, colorimeters, and electronic spectrometers. With a little practice, anyone can obtain accurate and reliable water quality data using these simple procedures.


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