Sulfur isotope biogeochemistry is the study of the distribution of sulfur stable isotopes in biological and geological materials. In addition to its common isotope, 32S, sulfur has three rare stable isotopes: 34S, 36S, and 33S. The distribution of these isotopes in the environment is controlled by many biochemical and physical processes, including biological metabolisms, mineral formation processes, and atmospheric chemistry. Measuring the abundance of sulfur stable isotopes in natural materials, like bacterial cultures, minerals, or seawater, can reveal information about these processes both in the modern environment and over Earth history.
Background
Natural abundance of sulfur isotopes
Sulfur has 24 known isotopes[1], 4 of which are stable (meaning that they do not undergo radioactive decay). 32S, the common isotope of sulfur, makes up 95.0% of the natural sulfur on Earth.[2] In the atomic symbol of 32S, the number 32 refers to the mass of each atom in Daltons, the result of the 16 protons and 16 neutrons of 1 Dalton each that make up the sulfur nucleus. The three rare stable isotopes of sulfur are 34S (4.2% of natural sulfur), 33S (0.75%), and 36S (0.015%) [3]. These isotopes differ from 32S in the number of neutrons in each atom, but not the number of protons or electrons. As a result, each isotope has a slightly different mass, but has nearly identical chemical properties.
Physical chemistry
Small differences in mass between stable isotopes of the same element can lead to a phenomenon called an "isotope effect," where heavier or lighter isotopes are preferentially incorporated into different natural materials. Isotope effects are divided into two main groups: kinetic isotope effects and equilibrium isotope effects. A kinetic isotope effect occurs when a reaction is irreversible, meaning that the reaction only proceeds in the direction from reactants to products. Kinetic isotope effects cause isotope fractionation--meaning that they affect isotope distributions--because the mass differences between stable isotopes can affect the rate of chemical reactions. It takes more energy to reach the transition state of a reaction if the compound has bonds with a heavier isotope, which causes the compound with heavier isotopes to react more slowly. Normal kinetic isotope effects cause the lighter isotope or isotopes to be preferentially included in a reaction's products. The products are then said to be "depleted" in the heavy isotope relative to the reactant. [4] Equilibrium isotope effects cause fractionation because it is more chemically favorable for heavy isotopes to take part in stronger bonds. [5] An equilibrium isotope effect occurs when a reaction is at equilibrium, meaning that the reaction is able to occur in both directions: simultaneously, reactants are forming products and products are returning to reactants. When a reaction is at equilibrium, heavy isotopes will preferentially accumulate where they can form the strongest bonds. For example, when the water in a sealed, half-full bottle is in equilibrium with the vapor above it, the heavier isotopes 2H and 18O will accumulate in the liquid, where they form stronger bonds, while the lighter isotopes 1H and 16O will accumulate in the vapor. The liquid is then said to be "enriched" in the heavy isotope relative to the vapor.
Calculations
Delta notation
Differences in the abundance of stable isotopes among natural materials are usually very small (natural differences in the ratio of rare to common isotope are almost always below 0.1%, and sometimes much smaller).[6] Nevertheless, these very small differences can record meaningful biological and geological processes. For this reason, isotope abundances in natural materials are often reported relative to isotope abundances in designated standards. By reporting all measurements relative to the same agreed-upon standard, small but meaningful differences in isotope abundances can be more easily understood and compared. The convention for reporting the measured difference between a sample and a standard is called "delta notation." For example, imagine an element X for which we wish to compare the rare, heavy stable isotope with atomic mass A (AX) to the light, common isotope with atomic mass B (BX). The abundance of AX and BX in any given material is reported with the notation δAX. δAX for the sample material is calculated as follows[7]:
AR = (total amount of AX)/(total amount of BX) δAXsample = (ARsample - ARstandard)/ARstandard
δ values are most commonly reported in parts per thousand, commonly referred to in isotope chemistry as per mille and represented by the symbol ‰. To report δ values in per mille, the δ value as calculated above should be multiplied by 1000:
While an isotope effect is the physical tendency for stable isotopes to distribute in a particular way, the isotopic fractionation is the measurable result of this tendency. The isotopic fractionation of a natural process can be calculated from measured isotope abundances. The calculated value is called a "fractionation factor," and allows the effect of different processes on isotope distributions to be mathematically compared. For example, imagine a chemical reaction Reactant → Product. Reactant and Product are materials that both contain the element X, and X has two stable isotopes, AX and BX. The fractionation factor for the element X in the reaction Reactant → Product is represented by the notation AαProduct/Reactant (although superscripts and subscripts are sometimes omitted). AαProduct/Reactant is calculated as follows[8]: AαProduct/Reactant = (δAXProduct - 1)/(δAXReactant - 1) Fractionation factors are often reported using the notation AεProduct/Reactant, calculated as follows: AεProduct/Reactant = AαProduct/Reactant - 1 Like δ values, ε values can be reported in per mille by multiplying by 1000.
Δ33S and Δ36S notation
All kinetic and equilibrium isotope effects result from differences in atomic mass. As a result, a reaction that fractionates 34S will also fractionate 33S and 36S, and the fractionation factor for each isotope will be mathematically proportional to its mass. Because of the mathematical relationships of their masses, the observed relationships between δ34S, δ33S, and δ36S in most natural materials are approximately δ33S = 0.515 × δ34S and δ36S = 1.90 × δ34S.[9] Rarely, natural processes can create deviations from this relationship, and these deviations are reported as Δ33S and Δ36S values, usually pronounced as "cap delta." These values are often calculated as follows:[10] Δ33S = δ33S - 0.515 × δ34S Δ36S = δ36S - 1.89 × δ34S However, Δ33S and Δ36S calculations are not standardized and can vary among publications.[11]
Reference materials
Agreed-upon reporting standards are required so that reported δ values are comparable among studies. For the sulfur isotope system, δ34S values are reported on the Vienna-Cañon Diablo Troilite (VCDT) scale. The original CDT scale was based on a sample of the mineral troilite recovered from the Canyon Diablo meteorite at Meteor Crater, Arizona, US. The Cañon Diablo Troilite was established as having a δ34S value of 0‰. However, troilite from the Canyon Diablo meteorite was later discovered to have variable sulfur isotope composition.[12] As a result, VCDT was established as a hypothetical reference material with a 34R value of 0.044151[13] and δ34S of 0‰, but no physical sample of VCDT exists. Samples are now measured in comparison to IAEA reference materials, which are well-characterized, lab-prepared compounds with known δ34S values. A commonly-used IAEA reference material is IAEA-S-1, a silver sulfide reference material with a δ34S value of -0.30‰ VCDT [14][15]. 33S and 36S abundance can also be measured relative to IAEA reference materials and reported on the VCDT scale. For these isotopes, too, VCDT is established as having δ33S and δ36S values of 0‰.[16] The 33R value of VCDT is 0.007877 and the 36R value is 0.0002. IAEA-S-1 has a 33R value of 0.0007878 and a δ33S value of -0.05‰ VCDT; it has a δ36S value of -0.6‰ VCDT.[17]A Canyon Diablo meteorite sample. The original reference standard for measuring δ34S was the mineral troilite (FeS) recovered from the Canyon Diablo meteorite
Analytical methods and instrumentation
The sulfur isotopic composition of natural samples can be determined by Elemental Analysis-Isotope Ratio Mass Spectrometry (EA-IRMS), by Dual Inlet-Isotope Ratio Mass Spectrometry (DI-IRMS), or by Multi-Collector-Inductively Coupled Plasma Mass Spectrometry (MC-ICPMS). MC-ICPMS can also be paired with gas chromatography (GC-MC-ICPMS) to measure some volatile compounds.
Natural variations in sulfur isotope abundance
Sulfur cycle background
Sulfur is present in the environment in solids, gases, and aqueous species. Sulfur-containing solids on Earth include the common minerals pyrite (FeS2), galena (PbS), and gypsum (CaSO4•2H2O). Sulfur is also an important component of biological material, including in the essential amino acids cysteine and methionine, the B vitamins thiamine and biotin, and the ubiquitous substrate coenzyme A. In the ocean and other natural waters, sulfur is abundant as dissolved sulfate. Hydrogen sulfide is also present in some parts of the deep ocean where it is released from hydrothermal vents.[18] Both sulfate and sulfide can be used by specialized microbes to obtain energy or to grow.[19] Gases including sulfur dioxide and carbonyl sulfide make up the atmospheric component of the sulfur cycle. Any processes that transport or chemically transform sulfur between these many natural materials also have the potential to fractionate sulfur isotopes.
Sulfur isotopic abundance in natural materials
Sulfur in natural materials can vary widely in isotopic composition. Compilations of the δ34S values of natural sulfur-containing materials include values ranging from -55‰ to 135‰ VCDT.[20] The ranges of δ34S values vary among sulfur-containing materials: for example, the sulfur in animal tissue ranges from ~ -10 to +20‰ VCDT, while the sulfate in natural waters ranges from ~ -20 to +135‰ VCDT.[21] The range of sulfur isotope abundances in different natural materials results from the fractionation processes that can affect those materials, discussed in the next section.
Known fractionation processes
Numerous natural processes are capable of fractionating sulfur isotopes. Important sulfur-fractionating processes include microbial metabolisms, mineral formation, high-temperature alteration of rock, and atmospheric chemical reactions. Microbes are capable of a wide variety of sulfur metabolisms, including the oxidation, reduction, and disproportionation (or simultaneous oxidation and reduction) of sulfur compounds. The effect of these metabolisms on sulfur isotopes in the reactants and products is also highly variable, depending on the rate of relevant reactions, availability of nutrients, and other biological and environmental parameters.[22][23] As an example, the microbial reduction of sulfate to sulfide generally results in a 34S-depleted product, but the strength of this fractionation has been shown to range from 0 to 65.6‰ VCDT.[24][25] Small fractionations with ε values from 0-5‰ have been observed in the formation of the mineral gypsum, an evaporite mineral produced through the evaporation of seawater.[26] Some sulfide minerals, including pyrite and galena, can form through thermochemical sulfate reduction, a process in which seawater sulfate trapped in rock is reduced to sulfide by geological heat; this process generally fractionates sulfur more strongly than gypsum formation.[27]Pyrite, a sulfur-bearing mineral that forms in some ocean sediments, usually has relatively low δ34S values due to the indirect role of biology in its formation.Prior to the rise of oxygen in Earth's atmosphere (referred to as the Great Oxidation Event), additional sulfur-fractionating processes referred to as mass-anomalous or mass-independent fractionation uniquely affected the abundance of 33S and 36S in the rock record. Mass-anomalous fractionations are rare, but they can occur through certain photochemical reactions between gases in the atmosphere. Studies have shown that mass-anomalous fractionation of sulfur isotopes could only have been recorded in ancient rock under conditions of very low atmospheric oxygen.[28]
Observed 34ε values for some common natural processes.
Signatures of mass-anomalous sulfur isotope fractionation preserved in the rock record have been used to infer the rise of atmospheric oxygen.[59] Nonzero values of Δ33S and Δ36S are present in the sulfur-bearing minerals of Precambrian rock formed greater than 2.45 billion years ago, but completely absent from rock less than 2.09 billion years old.[60] These fingerprints of mass-anomalous fractionation could only have been created and preserved if oxygen was essentially absent from Earth's atmosphere.[61][62] Δ33S and Δ36S measurements have been an important piece of evidence for understanding the Great Oxidation Event, the sudden rise of oxygen on Earth 2.45 billion years ago.[63]
Paleobiology and paleoclimate
A number of microbial metabolisms fractionate sulfur isotopes in distinctive ways, and the sulfur isotopic fingerprints of these metabolisms can be preserved in minerals and ancient organic matter. By measuring the sulfur isotopic composition of these preserved materials, scientists can reconstruct ancient biological processes and the environments where they occurred. δ34S values in the geologic record have been inferred to reveal the history of microbial sulfate reduction[64][65] and sulfide oxidation.[66] Paired δ34S and Δ33S records have also been used to show ancient microbial sulfur disproportionation.[67][68]
Microbial dissimilatory sulfate reduction (MSR), an energy-yielding metabolism performed by bacteria in anoxic environments, is associated with an especially large fractionation factor. The observed 34εMSR values range from 0 to -65.6 permil (see table for references). Many factors influence the size of this fractionation, including sulfate reduction rate,[69][70]sulfate concentration and transport,[71][72] availability of electron donors and other nutrients,[73][74][75] and physiological differences like protein expression.[76] Sulfide produced through MSR may then go on to form the mineral pyrite, preserving the 34S-depleted fingerprint of MSR in sedimentary rocks.[77] Many studies have investigated the δ34S values of ancient pyrite in order to understand past biological and environmental conditions. For example, pyrite δ34S records have been used to reconstruct shifts in primary productivity levels,[78] changing ocean oxygen content,[79][80] and glacial-interglacial changes in sea level and weathering.[81] Some studies compare sulfur isotopes in pyrite to a second sulfur-containing material, like sulfate or preserved organic matter.[82][83] Comparing pyrite to another material gives a fuller picture of how sulfur moved through ancient environments: it provides clues about the size of ancient 34εMSR values and the environmental conditions controlling MSR fractionation of sulfur isotopes.
Paleoceanography
δ34S records have been used to infer changes in seawater sulfate concentrations. Because the δ34S values of carbonate-associated sulfate are thought to be sensitive to seawater sulfate levels, these measurements have been used to reconstruct the history of seawater sulfate.[84] δ34S values of pyrite have also been applied to reconstruct the concentration of seawater sulfate, based on expected biological fractionations at low sulfate concentrations.[85][86] Both of these methods rely on assumptions about the depositional environment or the biological community, creating some uncertainty in the resulting reconstructions.[87][88]
Sulfur isotope biogeochemistry is the study of the distribution of stable sulfur isotopes in biological and geological materials. In addition to its common isotope, 32S, sulfur has three rare stable isotopes: 34S, 36S, and 33S. Measuring the abundance of these stable isotopes in natural materials can reveal information about biology and geology in the present and past.
Background
Natural abundance of sulfur isotopes
Sulfur has 24 known isotopes[89], 4 of which are stable (meaning that they do not undergo radioactive decay). 32S, the common isotope of sulfur, makes up 94.99% of the natural sulfur on Earth.[90]...describe natural abundance of other S isotopes, detail how they differ in number of neutrons and mass, but not in protons, electrons, chemical properties
Physical chemistry
Explain the basics of Isotope fractionation--preferential inclusion or exclusion of heavy isotopes in different materials because of small differences in mass
Coupling an elemental analyzer, using combustion analysis to quantitatively convert sample sulfur to SO2, to an isotope-ratio mass spectrometer able to measure the ratio of the stable sulfur isotopes in the samples
Over time, this technique has been improved to analyze increasingly small samples
Can be used for bulk sulfur-containing material, like sediment, biomass, etc.
Sulfur metabolisms: sulfate reduction, sulfur disproportionation as discussed in e.g. Johnston et al 2005 [93]
Effects of other metabolisms including oxidation, sulfite reduction, and diffusion effects as discussed in Chambers and Trudinger 2009 [94]
Discussion of what materials record the isotopic fingerprints of these metabolisms (e.g. pyrite vs. sulfate, etc.)
Mass-anomalous fractionations via photochemistry [95]
Applications
History of atmospheric oxygen
Using mass-anomalous fractionations to constrain the rise of atmospheric oxygen, e.g. Farquhar 2000 [96] and Pavlov and Kasting, 2002 [97]
Paleobiology
Using d34S values to understand ancient microbial activity, e.g. Bontognali 2012 [98], and changes in microbial metabolisms over time, e.g. Canfield and Teske 1996 [99], Shen 2009 [100], Shen and Buick 2004 [101]
Paleoceanography
Reconstructing seawater chemistry in the past, including sulfate and oxygen levels, e.g. Fike 2015 [102] and Habicht 2002 [103]
Environmental changes
Using pyrite, sulfate, and organic sulfur isotopes to understand major environmental changes, e.g. Raven 2019 [104] Pasquier 2017 [105], Fike 2018 [106], Owens 2013 [107] etc
Glacial-interglacial changes, ocean anoxia, extinctions, etc.
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