Isotopic fractionation

Processes as fundamental as evaporation, condensation, fluid movement through a rock, and many biological functions result in isotope fractionation. The record of these processes as they influenced the formation of various minerals is preserved in the distribution of stable isotopes within rock.

Weight Difference

Fractionation implies the breaking of a whole into its parts. In the case of a chemical element, the parts are the naturally occurring isotopes of that element. Some of the isotopes may be radioactive and may spontaneously decay to form another element. Most isotopes, however, are stable; they do not decay, and they differ from other isotopes only in their mass. Thus, stable isotope fractionation comprises several physical and chemical processes that can separate the stable isotopes of an element on the basis of weight difference.

The weight difference can be substantial. For example, deuterium is a stable isotope of hydrogen. The hydrogen nucleus contains one proton, whereas the deuterium nucleus contains one proton and one neutron. Because the proton and the neutron are about equal in mass (and the electron is so small it can be ignored), the atomic weight of hydrogen is 1 and the atomic weight of deuterium is 2; deuterium weighs 100 percent more than does hydrogen. In another example, the most common isotopes of carbon are carbon-12 (six protons and six neutrons in the nucleus) and carbon-13 (six protons and seven neutrons in the nucleus). Just as deuterium has one more neutron than does hydrogen, so carbon-13 has one more neutron than does carbon-12. The weight difference between carbon-13 and carbon-12, however, is less on a percentage basis—only 8 percent. It is apparent, then, that the addition of a neutron in an isotope has the greatest relative mass impact for the light elements (for example, hydrogen) and that the impact decreases as the elements become heavier. Thus, stable isotope fractionation processes are most obvious when light elements (those preceding sulfur in the periodic table) are involved.

Evaporation and Condensation

Fractionation includes those processes that separate light and heavy isotopes by physical means or through some chemical reaction. Evaporation and condensation are two physical processes that result in the separation of stable isotopes. For their impact on the earth's climate, the evaporation and condensation of water are arguably the most important physical processes.

Evaporation and condensation are mirror images of each other as far as isotope fractionation is concerned. In evaporation, energy is absorbed by water. This energy absorption is reflected by an increase in the temperature of the water. The individual water molecules absorb the energy and begin to move and to vibrate faster. Eventually, the individual molecules absorb so much energy that large quantities of them change from water molecules in a liquid to water molecules in a gas. The water molecules are said to have undergone a change of phase, from a liquid to a gas. If the process is reversed and energy is removed from the gas containing water molecules, the molecules slow down and begin to clump together. In the atmosphere, this condensation creates water droplets that may eventually produce a rain shower.

Water molecules are actually not all the same. With its two hydrogen atoms and one oxygen atom, the water molecule may have any of a range of molecular weights, depending upon the range of isotopic substitution. Stable isotopes of hydrogen (hydrogen and deuterium) and stable isotopes of oxygen (oxygen-16, oxygen-17, and oxygen-18) combine to yield water molecules of varying molecular weights. Normal water has a molecular weight of 18 (two hydrogens plus one oxygen-16), and heavier water has a molecular weight of 20 (two hydrogens and one oxygen-18). The heavier water molecule is 11 percent heavier than is light water.

If a mixture of light- and heavy-oxygen water is undergoing evaporation, it will take more energy input to “lift” the heavy water out of the liquid phase and into the gas phase. This process is analogous to a weightlifter's expending more energy to raise a 200-kilogram weight than a 180-kilogram weight. In the case of evaporation, the input energy is thermal (heat) rather than mechanical (physical lifting). The lighter oxygen-16 water will evaporate more readily, leaving the heavier oxygen-18 water behind. The gas phase above the evaporating water (for example, the atmosphere above the ocean) is dominated by light water, whereas the water remaining in the liquid phase is dominated by heavier water. This is not to say that heavy water molecules never evaporate; they do. Yet if one defines a number, R, as the ratio of heavy oxygen to light oxygen and measures R before the evaporation starts and then in the water vapor and the liquid water after some period of evaporation, the value of R will change. The ratio will be larger in the liquid (concentrating heavy oxygen) and smaller in the water vapor (dominated by light oxygen).

Magnitude of Isotopic Fractionation

Geochemists who study isotope fractionation use an equation which uses the R values calculated above to express the magnitude of isotopic fractionation. The equation δ = [(R(_sample) − R(_standard) ÷ R(_standard) ] × 1,000 defines a quantity called delta (δ), which represents the difference in the ratio of heavy oxygen to light oxygen between a sample (for example, the water vapor) and a standard (for example, seawater). This difference is multiplied by 1,000 and expressed as per mil (instead of percent).

The value of δ may be positive or negative, depending on the sign of the numerator. If the numerator is negative, R_sample is less than R_standard. In the evaporation example, that means that the sample has less heavy oxygen (or more light oxygen) than does the standard. The sample is depleted in the heavy isotope of oxygen and enriched in the light isotope of oxygen. In evaporation, the vapor phase is depleted in heavy oxygen and its δ value is negative. Conversely, the δ value for the remaining liquid phase will be positive (R_sample greater than R_standard), because the liquid phase is enriched in the heavy isotope of oxygen. When water vapor condenses to form raindrops, the first droplets to form are enriched in heavy-oxygen water (δ greater than 0), and the remaining vapor is enriched in light-oxygen water (δ less than 0).

Seawater as Standard

Seawater makes up greater than 90 percent of the liquid water at the earth's surface and makes a good standard for identifying fractionation effects. It should be apparent from the previous example that rainwater will be isotopically lighter (in both hydrogen isotopes and oxygen isotopes) than seawater. Similarly, evaporative waters, such as the waters of the Red Sea, will be isotopically heavier than is normal seawater.

If an oxygen-containing mineral forms in contact with water, some of the oxygen in the mineral will probably come from the water. If the water is isotopically light (from a freshwater, terrestrial environment, for example), the mineral oxygen will be relatively light as well. If the mineral forms in contact with an evaporative brine, the mineral oxygen should be enriched in the heavy isotope of oxygen. Thus, the stable isotopic composition of a mineral can tell scientists about the composition of ancient waters.

Bond Breaking

Chemical and biological reactions involve the breaking of chemical bonds. Just as it takes energy to evaporate a molecule of water, so it takes energy to break the bond between hydrogen and oxygen atoms in water or between carbon and oxygen atoms in carbon dioxide. The bond between atoms can be visualized as a spring. Inputs of energy cause the bond (spring) to vibrate. The atoms at the ends of the bond move apart as the bond stretches, and they move together as the bond restores the molecule to its original shape. If sufficient energy is applied, the atoms move so far apart (stretching the “spring”) that there is no force remaining to pull them back together. At this point, the bond has been broken.

Bond breaking occurs in any chemical reaction, whether it involves living organisms (a biochemical reaction) or proceeds without biological intervention (an inorganic reaction). In the case of inorganic reactions, the energy to break bonds is usually supplied by the environment in the form of heat. At low temperatures, such as normal room temperature, bonds involving heavy isotopes are less likely to be broken than are those involving only light isotopes. For the same amount of energy input, a spring with heavy weights on each end vibrates and stretches more slowly than does the same spring with lighter weights on the ends. This analogy shows why it takes more energy input to break bonds involving heavy isotopes. As the environmental temperature increases, however, the energy necessary to break bonds becomes readily available, and the degree of isotopic fractionation decreases. The slight differences in bond strength are insignificant at higher temperatures.

When living organisms are involved, the fractionation can be exaggerated. In biochemical reactions, the organism is often the source of energy for bond breaking. When the isotopes are chemically identical, it does not make sense for an organism to expend the extra energy to break heavy-isotope bonds when the same reaction with light isotopes uses less energy. Even in the case of photosynthesis, where sunlight is the primary energy source used to break the carbon-oxygen bond in carbon dioxide, isotopically light carbon dioxide is more likely to be photosynthesized than is isotopically heavy carbon dioxide. The result is that light isotopes are concentrated in the reaction products of biochemical reaction. Thus, biochemical molecules, which contain elements such as hydrogen, carbon, oxygen, nitrogen, sulfur, and phosphorus, tend to concentrate the lighter isotopes of those elements. The surrounding environment, be it lakewater, seawater, or sediments, tends to accumulate the heavier isotopes.

The mass spectrometer made it possible for scientists to identify and study the phenomenon of isotope fractionation. As its name implies, the instrument analyzes a spectrum based on mass, or weight. The rainbow is a common example of the spectrum of white light. The original light has been broken into its colorful components by passing through water droplets in the atmosphere that act like a glass prism. Each individual color represents light of a different wavelength. Water droplets and prisms, then, serve as simple wavelength spectrometers. In a mass spectrometer, a sample is analyzed for the range of masses of the elements it contains.

In an actual analysis of a geologic sample—coal, for example—the sample is converted through chemical processes into gases suitable for analysis in the mass spectrometer. The gases produced by the sample—in this case probably carbon dioxide, water, and some nitrogen and sulfur gases—are separated and purified by passing them through a series of freezing and drying steps. When a pure gas sample is obtained, that portion of the sample is injected into the mass spectrometer. Inside the mass spectrometer, the sample first enters an ionization chamber, where one or two electrons are stripped from the molecule. The molecule is now a positively charged ion (it has lost one or more negative electrons). The ionized gas sample next approaches a negatively charged metal plate with a hole in the middle. Because the sample gas is positively charged and the metal plate is negatively charged, the sample molecules are attracted to the plate and accelerate toward it. Some of the sample molecules pass through the hole.

Isotopic Separation

At this point, isotopic separation becomes important. For example, in nitrogen gas (composed of two nitrogen atoms), the nitrogen molecule may contain either nitrogen-14 or nitrogen-15 isotopes. The molecule may weigh 28, 29, or 30 units. During the acceleration toward the metal plate, the lighter-isotope molecules will be moving faster than the heavier molecules. Nitrogen gas molecules composed of two nitrogen 15 atoms are moving most slowly.

After passing through the hole, the nitrogen molecules enter a metal pipe with a slight bend. There is a vacuum in the pipe, and the pipe is surrounded by a strong magnet. As the accelerated, charged molecules enter the bent pipe, their flight path is bent by the presence of the magnetic field. The degree of bending depends on how fast the molecule is moving, which, in turn, depends on the mass of the molecule and the amount of charge on the metal accelerating plate. Thus, given one set of conditions (accelerating plate voltage and strength of the magnetic field), only the nitrogen gas molecules weighing 29 units will negotiate the bend in the pipe and reach the particle detector at the other end. On the one hand, the lighter nitrogen will be moving too fast to make the turn and will adhere to the outer wall of the pipe. On the other hand, the heavier nitrogen molecules will be moving too slowly and will be bent sharply into the inside wall of the pipe. By varying the accelerating voltage, the scientist can focus beams of all three types of nitrogen gas onto the detector and analyze the nitrogen isotopic composition of the sample. The same sort of analysis is done with hydrogen isotopes, using hydrogen gas and carbon, and with oxygen isotopes, using carbon dioxide gas.

Understanding Climate and Biochemical Processes

In general, a scientist analyzes a sample and seeks a measure of isotopic fractionation to determine either at what temperature a mineral formed or whether a mineral formed at lower temperatures is inorganic or biochemically influenced. The samples analyzed may occur naturally, or they may be minerals grown in carefully controlled laboratory systems where the precise isotopic compositions of the starting materials are known.

Stable isotope fractionation is a process that occurs continuously. The processes involved are as simple as evaporation and as fundamental to human survival as are photosynthesis and respiration. Many of the natural fractionation processes studied with stable isotopes are very simple, but others are extremely complex. Among the former, precipitation in the form of snow leaves a permanent, frozen record of climate in the polar ice caps. Studies of climatic change in these polar latitudes involve drilling ice cores out of the thick ice sheet and determining the stable isotope ratios in the water. Because human activities such as the burning of fossil fuels and forest clear-cutting are so extensive, some scientists believe that the earth's surface climate is being changed. Understanding the range of climate variability over the recent geologic past through the study of isotopic variations in ice cores may tell scientists whether the earth's climate can absorb and recover from such changes.

Biological processes would fit into the category of extremely complex processes. Scientists who study the details of biochemical processes must examine and understand each step in a series of many that make up the overall reaction. At any point in that series of reactions, if something goes wrong, the chance for the development of a disease or abnormality exists. Isotopes in general, and stable isotopes in particular, can be used to trace the starting chemicals through the maze of reactions to the ultimate products. Stable isotopes act to identify certain elements throughout the reaction sequence, without introducing possible radiation effects (as when radioactive isotopes are used). Knowledge of fractionation processes allows the scientist to correct for or ignore certain amounts of enrichment or depletion in the reaction products. Identifying stable isotopes as parts of nonproductive side branches of the reaction series may provide a clue to the reaction process, which, in turn, may suggest a weakness in the process. Understanding important biochemical reactions on the molecular scale may help scientists to uncover cures for diseases and prevent birth defects.

Principal Terms

chemical bond: the force holding two chemical elements together as part of a molecule

depletion: the process by which the light isotope is concentrated in either the reactants or the products of a chemical reaction

enrichment: the process by which the heavy isotope is concentrated in either the reactants or the products of a chemical reaction

isotopes: atoms containing the same number of protons but a different number of neutrons, giving the same chemical properties but different atomic weights

isotopic fractionation: changes in the isotopic composition of natural substances, which result from small differences in the physical, chemical, and biological properties of isotopes

product: the material that results from a reactant undergoing a chemical process

reactant: the starting material or materials in any chemical reaction

standard: a material of known isotopic composition; all enrichment and depletion is measured relative to the standard value

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