Freshwater chemistry

The unique properties of water that are so important to geological processes depend upon the distinctive polar structure of the water molecule. The chemistry of freshwater is highly variable. Throughout the earth, it is influenced by the atmosphere and locally influenced by reactions with the local soils and bedrocks. Changes through time in the chemistry of natural freshwater is one of the best indicators of changes in the earth's surface environment.

Polar Molecule

The temperature of the earth lies within the range that permits water to exist as a liquid, a gas, and a solid. Earth is sometimes called the “blue planet” because of the brilliant blue color, seen from outer space, that results from more than three-quarters of the planet's surface being covered by water. Ultimately, this water comes from within the planet during magmatic differentiation and formation of the crust. The size of the earth provides a gravitational field that is sufficient to keep water vapor from escaping through geologic time. Steam, the gaseous form of water, is the dominant constituent of volcanic gases, and pressure generated by steam accounts for the power of most explosive eruptions. Ice in glaciers has been a major force in influencing the topography of vast areas of the continent.

The chemical symbol for water, H₂O, refers to a molecule composed of one hydrogen ion (H⁺) and one hydroxyl ion (OH⁻). When combined to form a water molecule, the two hydrogen ions are arranged to one side and have an angle of about 105 degrees between them. Because the positively charged hydrogen ions lie on one side of the molecule, they impart a positive charge to that side of the molecule while leaving a negative charge of the oxygen exposed at the other end. Therefore, each water molecule is like a tiny magnet, with a positive pole and a negative pole. The molecule has polar charge distributions and thus is termed a polar molecule.

Because each water molecule is like a magnet, as a substance, water will behave somewhat like a box full of very tiny magnets. The tendency for the magnets would be to line up positive pole to negative pole, and a similar tendency occurs in water, where the ionic attraction of the hydrogen of one water molecule for the oxygen of another water molecule actually draws water molecule to water molecule and creates a type of ordering. The cohesive ordering that results from this attraction is called hydrogen bonding, and it accounts for some remarkable physical properties that belong only to water.

Water is one of the few liquid substances that do not simply become denser as they cool. The hydrogen bond allows liquid water molecules to pack themselves into a tighter pattern than would be possible if the molecules were not polar. Water achieves its maximum density of 1.00000 gram per cubic centimeter at 3.94 degrees Celsius. Liquid water molecules vibrate rapidly (about 10¹² vibrations per second), but when the water is cooled to the freezing temperature of ice (0 degrees Celsius), these vibrations decrease to about 10⁶ vibrations per second. At that point, new covalent bonds are able to overcome the force of the hydrogen bond. The slowing of vibrations as water cools permits some rearrangement and some creation of open space to begin at just below 3.94 degrees Celsius, even though ice crystals do not form until 0 degrees is reached. This accounts for the maximum density of water occurring at 3.94 degrees rather than at 0 degrees Celsius.

When actual freezing occurs, water molecules arrange themselves into a covalently bonded crystal structure that has open space not present in liquid water. Water goes from a density of 0.99987 gram per cubic centimeter as liquid to 0.917 gram per cubic centimeter as solid ice. The pronounced decrease in density causes ice cover to form on the surface of lakes and rivers rather than at the bottom. Without this property, aquatic life would probably not be possible, since ice on the surface insulates the water below, keeping it above 0 degrees Celsius and preventing complete freezing from bottom to top. The decrease in density is accompanied by an increase in volume; water expands as it freezes. When water seeps into cracks and pores of rock and soil and then freezes, the force generated by the expansion causes rock to break, soils to heave and swell, and small grains and crystals to spall away. Therefore, it is the hydrogen bond of water that ultimately permits water to become a powerful mechanical agent in weathering, as exemplified through frost heave in soils or frost wedging in rock.

Above 3.94 degrees Celsius, the density of water decreases with increasing temperature. At temperatures likely to be encountered in lakes and streams, the density of pure water can be closely approximated with a formula. The small changes in density that occur during heating, cooling, and freezing are responsible for seasonal circulation of water in freshwater lakes.

Water is unlike other substances in that its freezing point is lowered, rather than raised, by pressure, a result of a tendency at near-freezing temperatures for the covalent bonds of ice to collapse back to the hydrogen bonds of liquid water. Simple pressure is sufficient to enact the transition. When an ice skater places pressure on the blades of the skate, the pressure melts the ice at the sharp edges of the blade, and the skater glides on a thin layer of water that is created momentarily by the pressure. Such skating is not possible on other substances such as solid carbon dioxide (dry ice). This same liquefaction under pressure permits glacial ice to flow over a thin layer of water created at the base of the glacier and for this water to flow into joints and cracks, refreeze, and “pluck” out sections of the bedrock as the glacier moves.

Electrostatic Bonding

Water is a powerful solvent because the water molecule is polar. The magnet-like quality of the water molecule strongly influences substances that are bonded primarily by electrostatic (ionic) bonds. Ions bound into many solid minerals by virtue of the ionic bond may therefore be dissolved into water. In this way, water picks up dissolved materials from the soils and rock. The dissolving power provided by the polar nature of the molecules is further assisted by their vibration. As the molecules vibrate faster, they act like high-speed jackhammers against solid particles, and those solids held together with ionic bonds are particularly affected. Heating is one way to increase the vibration of molecules. Therefore, freshwater tends to be a more effective solvent in warm climates.

Freshwater contains dissolved solids and gases. The substances that are likely to occur in waters of this planet are those substances that are abundant and that occur in forms that are easily soluble. More than 99 percent of the earth's crust can be accounted for by only twelve elements (oxygen, silicon, aluminum, iron, calcium, magnesium, sodium, potassium, titanium, hydrogen, manganese, and phosphorus—in order of abundance from greatest to least), and nearly 100 percent of the atmosphere by five elements (nitrogen, oxygen, argon, hydrogen as water vapor, and carbon as carbon dioxide).

The most common dissolved substances in freshwater are bicarbonate, calcium, sodium, magnesium, potassium, fluorine, iron, phosphate, sulfate, chloride, and dissolved silica. These substances enter natural water both from the atmosphere and from the rocks and soils contacted by the water. Even rainwater is not completely pure and contains small amounts of silica, sulfate, calcium, chlorine, nitrate, carbonic acid, sodium, potassium, iron, and aluminum.

Some elements are crustally abundant but are not abundant as dissolved species in freshwater because they form insoluble compounds. For example, aluminum and iron are abundant elements, but in most water they form nearly insoluble hydroxides—aluminum hydroxide and iron hydroxide. Silicon is extremely abundant, but it forms strong covalent bonds with oxygen that are not possible for water to break. In nature, silicon forms minerals that include silica-oxygen structures with very low solubility in water.

Acidic and Alkaline Waters

Although water is a good solvent because of the polar nature of the water molecule, it can be made an even more effective solvent if it can be rendered acidic. Chemists have devised the pH scale as a means for expressing the degree to which a solution is acid. The scale runs from 0 (most acid) to 14 (least acid), with a pH value of 7 termed neutral. Waters with a pH of less than 7 are termed acidic, and those with pH values greater than 7 are termed alkaline. Natural rainwater is acidic and has a pH of 5.6, because carbon dioxide in the atmosphere dissolves into the rainwater to make carbonic acid. (The atmosphere has a powerful effect on the composition of freshwater; the character of freshwater on this planet has changed through geologic time in accord with changes in the atmosphere.) Once rainwater strikes the ground, the weak acid is neutralized as it dissolves minerals in the rocks and soils. Therefore, even though natural rainwater is acidic, most fresh groundwater and surface water is near neutral, with pH values between 6 and 8.5.

Unusually acidic waters occur around volcanic vents, geysers, and fumaroles, where gases such as sulfur dioxide react with water to produce sulfuric acid. Acidic water also occurs in bogs and marshes (pH 3.3-4.5), where carbon dioxide released during the decay of organic matter reacts with water to produce carbonic acid. The strongest acids occur where sulfide minerals such as pyrite and marcasite oxidize in the presence of air and water to produce sulfuric acid. Sulfuric acid is a strong acid, and its production through the weathering of sulfide minerals is so common around coal and metal mines that the water released are termed acid mine drainage. Streams flowing from these sites often have pH values below 2. Thousands of miles of streams have become polluted by acid mine drainage. Alkaline water occur in limestone terrains, where pH values between 8 and 9 are common. Extremely alkaline water occurs in playa lakes in contact with sodium carbonate or sodium borate, where pH values above 12 have been noted.

Extremely acidic and extremely alkaline water permit substances that remain immobile in normal water to be dissolved. For example, aluminum is normally insoluble but does dissolve in both very acidic and very alkaline water. Acidic water dissolves and transport metals such as iron and manganese and nonmetals such as phosphorus that would normally be present in most natural water at low or undetectable levels.

Acidity vs. pH

A clear distinction needs to be made between pH and acidity, because these terms are too often misused interchangeably. In a large tank of pure distilled water at pH = 7, if a few drops of strong acid (a substance that releases free positively charged hydrogen ions into solution) are added, a low pH, perhaps pH = 3 throughout the tank, will soon be registered. If a few drops of strong base (a substance that releases free negatively charged hydroxide ions into solution) are added, however, the pH will easily rise back to 7. The measure of the amount of base required to get the tank to a pH of 7 is a measure of acidity. The tank had a low pH, but yet also a low acidity. Suppose another tank contains water with a considerable amount of dissolved solids and has a pH of 6. It may take many gallons of strong base to get the tank up to a pH of 7. In the second tank, then, there is water with a moderate pH but a very high acidity. In short, pH is the measure of the concentration of hydrogen ion present at a given time, but acidity is a measure of the amount of base needed to change the pH back to neutral. Natural bodies of water with high acidity are usually high in dissolved solids and have a low pH. Acid mine drainage is an example.

Aqueous Geochemistry

Freshwater chemistry is a part of the field of aqueous geochemistry, or low-temperature geochemistry, which is a specialty field of many geologists. Proper collection, analyses, and interpretation of data from natural water require more than knowledge of chemistry. Success also requires knowledge of the natural environmental system, and this is why hydrology and aqueous geochemistry fall more properly within the province of geology than that of pure chemistry.

A typical water analysis includes determination of the levels of silica, aluminum, iron, manganese, copper, calcium, magnesium, strontium, sodium, potassium, dissolved oxygen, carbonate, nitrate, bicarbonate, sulfate, chloride, fluoride, phosphate, arsenic, selenium, boron, total dissolved solids, pH, acidity or alkalinity, temperature, and conductivity. If water is polluted with unusual substances such as pesticides, sewage, or industrial wastes, specialized tests must be undertaken for these.

As soon as a sample is taken, the water is removed from its actual environment, and changes start to occur. Dissolved gases may leave, precipitates may form, and bacteria and microscopic algae may metabolize substances or die and release substances that were formerly solids. Because some parameters change so quickly and easily, some tests must be done immediately in the field. The temperature, conductance, pH, dissolved oxygen, and sometimes acidity or alkalinity is measured in the field. Small battery-operated instruments such as specific ion meters (for nitrate, sulfate, carbonate, and a number of elements), colorimeters, and conductivity meters are made by several manufacturers to be used in field analyses.

Pure water is a poor electrical conductor, but water's ability to conduct electricity increases markedly with the amount of dissolved solids. Therefore, a conductance test in the field is a rapid method to use to make a rough estimate of the amount of total dissolved solids. Sometimes rapid colorimetric tests for sulfate, phosphate, and some metals are done in the field, but more often the analyses of those substances are done in the laboratory.

Samples to be taken to the laboratory are filtered in the field to remove any microscopic suspended solids, because it is important that the water analyses show only dissolved substances. After the water is filtered, it is acidified with a few milliliters of high-purity acid to ensure that all dissolved species remain dissolved. The water is then put into a plastic container and filled to the top so as not to admit any air. The samples are usually placed in a dark cooler and kept refrigerated until they are analyzed in the laboratory.

Danger of Contaminants and Pollutants

Of the vast amount of water present on the planet's surface, only about 2.5 percent is freshwater; the remainder is in oceans and inland seas. Of that 2.5 percent, 2.15 percent is locked away in glaciers and polar ice caps. About another 0.3 percent of the total is accounted for by fresh groundwater. All the surface water commonly seen in freshwater lakes and streams amounts to only .009 percent of the earth's total water.

The small percentages do not reveal the true importance of freshwater. Freshwater is the major sculpting agent of the planet's land surface, and it is an essential substance for all terrestrial life-forms. Changes in the chemical composition of freshwater constitute one of the most sensitive indicators of changes in the environment. Changes may be local, as exemplified by the eutrophication of a small lake that receives phosphorus from a few local septic tanks, or they may be global, as exemplified by the acidification of sensitive lakes in northern Europe and Canada as a result of the burning of fossil fuels, which releases carbon dioxide and oxides of nitrogen and sulfur into the atmosphere.

The category “freshwater” includes water from wells, lakes, and streams with diverse chemistry. Even natural substances in excessive amounts pose problems for water consumption or industrial use. In addition to the toxic content of acid mine drainage, excessive amounts of nitrate, sodium, and fluoride present health hazards. Excessive amounts of calcium and magnesium promote the buildup of scale in water tanks and boilers, and water that is high in these elements (hard water) must have these constituents chemically removed with a water softener prior to domestic and industrial use. High phosphate content of water promotes algal blooms and eutrophication. High boron content is important in irrigation, because boron is toxic to many crops. Of all natural resources, freshwater ranks as one of the most essential to human survival. Continual monitoring of the chemistry of bodies of freshwater is essential to ensure that these valuable resources are not rendered unsuitable by human-made contaminants and pollutants.

Principal Terms

acid: a substance that yields free hydrogen ions in solution

acidity: the degree of a solution's being acidic as determined by the quantity of base needed to neutralize the solution

alkalinity: the degree of a solution's being basic as determined by the quantity of acid needed to neutralize the solution

anion group: a combination of ions that behaves as a single anion

colorimeter: an instrument that measures the intensity of color produced when a reagent reacts with a substance in a solution; the intensity of color is used to quantify the amount of the substance in solution

covalent bonding: a type of chemical bonding produced by sharing of electrons between overlapping orbitals of adjacent atoms; covalently bonded solids usually have low solubility in water

density: a property of a substance expressed in units of weight per unit of volume, such as pounds per cubic foot or grams per cubic centimeter

eutrophication: processes causing water bodies to receive excess nutrients that stimulate excessive plant growth

ionic bonding: a type of chemical bonding that holds the constituents of a crystal together primarily by electrostatic attraction between oppositely charged ions

total dissolved solids (TDS): a quantity of solids, expressed in weight percent, determined from the weight of dry residue left after evaporation of a known weight of water

Bibliography

Berner, Elizabeth K., and Robert A. Berner. Global Environment: Water, Air, and Geochemical Cycles. Upper Saddle River, N.J.: Prentice Hall, 1996.

Dodds, Walter K., and Matt R. Whiles. Freshwater Ecology: Concepts and Environmental Applications of Limnology. 2d ed. Burlington: Academic Press, 2010.

Drever, James I. The Geochemistry of Natural Waters. 3rd ed. Englewood Cliffs, N.J.: Prentice-Hall, 1997.

Faust, Aly. Chemistry of Natural Waters. Stoneham, Mass.: Butterworth Publishers, 1981.

Greenberg, Arnold, et al., eds. Standard Methods for the Examination of Water and Wastewater. 21st ed. Washington, D.C.: American Public Health Association, 2005.

Hem, J. D. Study and Interpretation of the Chemical Characteristics of Natural Water. US Geological Survey Water Supply Paper 2254. Honolulu: University Press of the Pacific, 2005.

Krauskopf, Konrad B. Introduction to Geochemistry. 3rd ed. New York: McGraw-Hill, 2003.

Lam, Buuan, et al. “Major Structural Components in Freshwater Dissolved Organic Matter.” Environmental Science & Technology 41 (2007): 8240-8247.

Langmuir, Donald. Aqueous Environmental Geochemistry. Upper Saddle River, N.J.: Prentice Hall, 1997.

Nicholson, Keith. Geothermal Fluids: Chemistry and Exploration Techniques. Berlin: Springer-Verlag, 1993.

Robinson, Philip. "There's a World of Chemistry in Water." Chemistry World, 25 Mar. 2024, www.chemistryworld.com/opinion/theres-a-world-of-chemistry-in-water/4019215.article. Accessed 9 Feb. 2025.

Sparks, Donald L., and Timothy J. Grundl, eds. Mineral-Water Interfacial Reactions: Kinetics and Mechanisms. Washington, D.C.: American Chemical Society, 1998.

Stumm, Werner, and James J. Morgan. Aquatic Chemistry: An Introduction Emphasizing Chemical Equilibria in Natural Waters. 3rd ed. New York: John Wiley & Sons, 1996.

van der Leeden, Fritz, Fred L. Troise, and D. K. Todd. The Water Encyclopedia. 2d ed. New York: CRC Press, LLC, 1990.