Isotopes: Applications in Natural Waters
Efforts to understand how natural waters evolve are aided significantly by isotope chemistry. The origin and type(s) of the water, the environment in which the water has resided, the types of reactions that the water has experienced, and the length of time the water has been removed from the hydrologic cycle are just some of the areas where isotopic chemistry has helped hydrologists understand natural waters.
A simplistic view of the atom is that it consists of a nucleus comprising protons and neutrons surrounded by electrons at varying distances and occupying orbitals of different shapes and orientations. Nuclear chemists and physicists have demonstrated that protons and neutrons are not indivisible in themselves, rather they consist of even smaller particles such as quarks, muons, gluons, and many others with exotic names and properties. For the purpose of this discussion, however, consideration of these sub-subatomic particles is not necessary to understand isotopes and their application to water chemistry.
Protons, positively charged particles within the nucleus, are a fundamental characteristic of atoms. The atomic number, or number of protons in the nucleus of an atom, controls the identity of the atom. If a nucleus contains 8 protons, that atom is oxygen. If the atom has 7 protons, it is nitrogen; 9 protons, it is fluorine. Changing the identity of the atom, the goal of alchemists in the Middle Ages, is accomplished on Earth only through radioactive decay or by the natural or artificial bombardment of a nucleus with high-speed particles.
The chemical behavior of an atom is controlled by the outermost electrons that surround the nucleus. Most elements can gain, lose, or share electrons. This ability controls which elements combine (bond) to form molecules, and in what proportions and orientation with respect to one another. Electrons have very small masses compared to protons, and possess a negative charge.
In addition to protons, the nucleus of an atom contains neutrons. Neutrons possess no charge and have a mass very similar to that of protons. Unlike protons, whose number controls an atom's identity, atoms of the same identity can have a different number of neutrons within the nucleus. The variation can occur only within specific limits, and has the impact of changing the mass of the particular atom. For example, oxygen with 8 protons in the nucleus can have 8, 9, or 10 neutrons. These different atoms, all with 8 protons in the nucleus, but having atomic masses of 16, 17 and 18 (i.e., 16 O, 17 O, and 18 O, respectively), are all isotopes of oxygen. The term isotope is applied to any atom in a set of atoms having the same atomic number, but differing in the number of neutrons (i.e., their mass).
Because the isotopes have the same chemical identity and outer electron configuration, they have similar chemical characteristics. The slightly differing mass, however, means that the different isotopes of an element will experience small differences in properties that influence reactions or reaction rates. For example, the different masses might mean differing bond energies, or may influence the rate at which the isotopes diffuse (spread out) through a substance, or in which they participate in a given reaction. Because of these small differences, and because technology has evolved to the state at which very small differences in the isotopic compositions can be measured (e.g., the proportions of 18 O, 17 O, and 16 O), scientists are able to use isotopes to identify specific reactions and their rates in natural waters.
Isotopic compositions are reported relative to a standard substance having a known ratio. Typically, the ratios, expressed as the heavy isotope to lighter isotope (e.g., 18 O/ 16 O) are reported as δ (delta) values where
δ = [Ratio in Sample − Ratio in Standard] / [Ratio in Standard] × 1000
Delta values can be positive or negative depending on whether the sample has a higher or lower isotopic ratio, respectively, than the standard (see table below).
|Hydrogen (H)||2 H / 1 H or D / H||Standard Mean Ocean Seawater|
|Oxygen (O)||18 O / 16 O||Standard Mean Ocean Seawater|
|Carbon (C)||13 C / 12 C||Fossil belemnite from a limestone||formation|
|Nitrogen (N)||15 N / 14 N||Atmosphere|
|Sulfur (S)||34 S / 32 S||Iron sulfide mineral in Canon||Diablo meteorite|
The isotopic composition of oxygen (δ 18 O)in a given sample, therefore, would be determined as follows:
δ 18 O sample = [( 18 O/ 16 O) sample − ( 18 O/ 16 O) seawater ] / [( 18 O/ 16 O) seawater ] × 1000
Application of Isotopes to Hydrology
Although isotopes occur naturally for virtually all elements, those commonly used in water-related studies are for hydrogen (H), carbon (C), nitrogen (N), oxygen (O) and sulfur (S). These elements are very common, and, because of their low total masses, the difference in the masses of the isotopes of a given element are proportionally large. As a result, small changes in the environment can produce a readily measurable difference in the ratios of the isotopes.
Isotopes are applied in several categories of hydrologic investigations. These include tracing water through the hydrologic cycle, paleothermometry , determining the origin of groundwater, the chemical environment of an aquifer , the age of groundwater, and the tracing and identification of pollutant sources. Commonly, the factors that are most important in determining the isotopic composition of a given element are the source (reservoir) of the element (e.g., atmosphere, sea water, rocks, organic matter), the characteristics of the environment (e.g., temperature, pH, redox conditions ), and the types of reactions that take place (e.g., photosynthesis , dissolution of rock, oxidizing reactions).
Oxygen and Hydrogen Isotopes and the Hydrologic Cycle.
Within the hydrologic cycle, evaporation and condensation reactions separate (fractionate) the heavy isotopes from the light isotopes of the elements H and O. Hydrogen has three naturally occurring isotopes, hydrogen (H), deuterium (D) and tritium (T), having 0, 1, and 2 neutrons in the nucleus, respectively. Oxygen has three common isotopes, 16 O, 17 O, and 18 O, as described above. When water evaporates, the lighter isotopes are preferentially enhanced in the vapor. During condensation of vapor to liquid, the reverse is true.
Origin of Groundwater.
The relationship between groundwater and meteoric water was firmly established when it was demonstrated that most waters, whether from a well, from a hot spring, or from a geyser , had hydrogen and oxygen isotopic compositions that could be related directly to the trend of meteoric water (established by Craig ). This observation confirmed that nearly all groundwater originates from the hydrologic cycle as precipitation. Whether a situation involves well water from Wisconsin, hot springs from Oregon, or geysers in Yellowstone National Park, in Alaska, or in New Zealand, the isotopic composition of the groundwater indicates that local precipitation is the origin.
Seasonal correlations between temperature and δ 18 O have been recognized in a number of places throughout the world, in part because temperature strongly influences evaporation and condensation. As a result, δ 18 O-enriched precipitation is found in the warmer areas. Also influencing the temperature–δ 18 O relationship are factors such as latitude, continental effects, and local characteristics such as altitude and seasonal variations.
The correlation between temperature and δ 18 O in precipitation is recorded in glacial ice. Ice cores show alternating layers reflecting seasonal atmospheric temperatures and record long-term temperature variations as well. This data allows scientists to reconstruct climatic conditions in the past.
Carbon Isotopes in Groundwater.
Carbon-13 (δ 13 C) provides a tool to determine the origin of carbon, either as dissolved inorganic or organic carbon. Carbon can originate from a number of sources (e.g., limestone, atmospheric CO 2 , plant organic matter). Limestones are important aquifers in the world. They consist of the mineral calcite (CaCO 3 ), which contains abundant inorganic carbon. Most of this carbon has an isotopic composition very close to that of the standard (see the listing at bottom of page 240), also a limestone. Dissolved inorganic carbon (DIC) from atmospheric CO 2 has an isotopic composition that differs from that of the DIC from limestone and is influenced by the pH of the solution.
During photosynthesis, carbon isotopes are fractionated depending on the type of plant and the photosynthetic cycle the plant follows. When these plants decay through respiration processes, they produce dissolved organic carbon (DOC) in the form of humic and fulvic acids. The δ 13 C of the DOC will reflect the specific plant origin.
As carbon, either as DIC and DOC interact, the resulting δ 13 C of the solution reflects the different carbon reservoirs (e.g., limestone, green plants), the conditions of the interaction (e.g., the pH and availability of CO 2 ), and the types of reactions involved (e.g., dissolving of calcite, photosynthesis, respiration).
Isotopes of Nitrogen.
Nitrogen is an element that is involved in many biological reactions. It is frequently released to the environment through the decay of organic matter. Nitrogen as nitrate is a risk to health, being linked to methemoglobinemia ("blue baby syndrome") in infants. Some research also points to the role of nitrate in some cancer occurrences.
Nitrate contamination commonly arises from two distinct sources: commercial fertilizers and human or animal wastes. The determination of which of these two sources of nitrate has detrimentally affected local groundwater may in some cases be resolved through the determination of the isotopic composition of nitrogen (δ 15 N) in the nitrate. Commercial fertilizer is made with atmospheric nitrogen and therefore has an isotopic composition of the standard. Organic wastes, on the other hand, lose nitrogen through ammonia (NH 3 ) loss, which preferentially takes the lighter of the nitrogen isotopes ( 14 N). As a result, nitrate derived from human waste (e.g., septic systems, animal waste from feedlots) has a significantly higher δ 15 N than that from commercial fertilizers. In actuality, the interpretation may be more difficult, owing to denitrification and the mixing of more than one source.
Carbon isotopes (from the DOC of the wastes) and oxygen isotopes (the oxygen in nitrate) are being used to further study nitrate contamination.
Isotopic compositions of many other elements commonly found within the Earth's waters provide important information that helps indicate how these waters have evolved. Examples of such "clues" include the sources of elements such as C, O, N, H, and S; the reactions that distribute them between different molecules; solid and liquid phases; and the environmental parameters temperature, pH, and redox conditions. Importantly, most of the elements above are independent of one another, providing for a means of testing and refining hypotheses of water evolution.
SEE ALSO C AVERN D EVELOPMENT ; F RESH W ATER , N ATURAL C OMPOSITION OF ; G LACIERS AND I CE S HEETS ; G ROUNDWATER ; G ROUNDWATER , A GE OF ; H YDROLOGIC C YCLE ; I CE A GES ; I CE C ORES AND A NCIENT C LIMATIC C ONDITIONS ; L ANDSLIDES ; R ADIONUCLIDES IN THE O CEAN ; S TREAM E ROSION AND L ANDSCAPE D EVELOPMENT ; T RACERS IN F RESH W ATER ; V OLCANOES AND W ATER ; W EATHERING OF R OCKS .
Dennis O. Nelson
Clark, Ian, and Peter Fritz. Environmental Isotopes in Hydrogeology. New York: Lewis Publishers, 1997.
Craig, Harmon. "Isotopic Variations in Meteoric Waters." Science. 133 (1961): 1702–1703.
Kendall, Carol, and Jeffrey J. McDonnell, eds. Isotope Tracers in Catchment Hydrology. New York: Elsevier, 1998.
MASS SPECTROMETRY: DETERMINING ISOTOPIC RATIOS
Small variations between isotope abundances of oxygen, hydrogen, carbon, nitrogen, and sulfur are measured using a machine called a gas-source mass spectrometer. The mass spectrometer measures the difference in the ratio of heavy to light isotope of a given element rather than measuring the actual abundances.
In the mass spectrometer source, the sample, converted to the gas H 2 , N 2 , CO 2 , or SO 2 , as appropriate, is heated and ionized (given a positive charge by removal of an electron—for example, CO 2 + ). In the high-vacuum source, the nowionized gas molecules are driven to high velocity across an electrical potential field (positive to negative charge), focused into a narrow beam, and directed into a flight tube (still under vacuum).
A strong magnetic field in the flight tube causes the path of the charged molecules to bend. The amount of the bending of the flight of a molecule depends on its mass and charge. As a result, molecules containing different isotopes (masses) of the element strike the collector at the end of the flight tube at different spots, allowing the ratio of the isotopic masses to be precisely measured.