Sulfur isotope biogeochemistry

Bar chart depicting the natural abundance of ³²S, ³³S, ³⁴S, and ³⁶S on Earth.

Sulfur has 24 known isotopes,[2] 4 of which are stable (meaning that they do not undergo radioactive decay).[3] 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 sulfur 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%).[4] 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.[3]

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 depending on the materials' chemical composition or physical state.[5] Isotope effects are divided into two main groups: kinetic isotope effects and equilibrium isotope effects.[5] A kinetic isotope effect occurs when a reaction is irreversible, meaning that the reaction only proceeds in the direction from reactants to products.[3][5] Kinetic isotope effects cause isotopic fractionation—meaning that they affect the isotopic composition of reactant and product compounds—because the mass differences between stable isotopes can affect the rate of chemical reactions.[5] 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.[5] Normal kinetic isotope effects cause the lighter isotope (or isotopes) to be preferentially included in a reaction's product.[5] The products are then said to be "depleted" in the heavy isotope relative to the reactant.[3] Rarely, inverse kinetic isotope effects may occur, where the heavier isotope is preferentially included in a reaction's product.[5][6]

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.[3] When a reaction is at equilibrium, heavy isotopes will preferentially accumulate where they can form the strongest bonds.[3] 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.[7] The liquid is then said to be "enriched" in the heavy isotope relative to the vapor.[3]

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).[5] Nevertheless, these very small differences can record meaningful biological and geological processes. To facilitate comparison of these small but meaningful differences, isotope abundances in natural materials are often reported relative to isotope abundances in designated standards.[3][5] 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 δ^AX. δ^AX for the sample material is calculated as follows:[5]

^AR = (total amount of ^AX)/(total amount of ^BX)

δ^AX_sample = (^AR_sample - ^AR_standard)/^AR_standard

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:

δ^AX_sample (‰) = ((^AR_sample - ^AR_standard)/^AR_standard) * 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.[5] 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.[5] 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 (the heavy isotope, with a mass of A) and ^BX (the light isotope, with a mass of B). The fractionation factor for the element X in the reaction Reactant → Product is represented by the notation ^Aα_{Product/Reactant}. ^Aα_{Product/Reactant} is calculated as follows:[5]

^Aα_{Product/Reactant} = (δ^AX_Product + 1)/(δ^AX_Reactant + 1)

Fractionation factors can also be reported using the notation ^Aε_{Product/Reactant}, which is sometimes called the "enrichment factor" and is calculated as follows:[5]

^Aε_{Product/Reactant} = ^Aα_{Product/Reactant} - 1

Like δ values, ε values can be reported in per mille by multiplying by 1000.

All kinetic and equilibrium isotope effects result from differences in atomic mass.[3][5] 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.[3] 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.[8] 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 typically calculated as follows:[3][9]

Δ³³S = 1000 × [(1 + δ³³S/1000) - (1 + δ³⁴S 1000)^0.518 - 1]

Δ³⁶S = 1000 × [(1 + δ³⁶S/1000) − (1 + δ³⁴S/1000)^1.91 − 1]

However, the method for calculating Δ³³S and Δ³⁶S values is not standardized, and can differ among publications.[10]

A Canyon Diablo meteorite sample. The original reference standard for measuring δ³⁴S was the mineral troilite (FeS) recovered from the Canyon Diablo meteorite.

Agreed-upon reference materials 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.[11] The original CDT scale was based on a sample of the mineral troilite recovered from the Canyon Diablo meteorite at Meteor Crater, Arizona, US.[3] The Cañon Diablo Troilite was assigned a δ³⁴S value of 0‰.[3] 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 sulfur isotope reference with a ³⁴R value of 0.044151[11] and δ³⁴S of 0‰, but no physical sample of VCDT exists. Samples are now measured in comparison to International Atomic Energy Agency (IAEA) reference materials, which are well-characterized, lab-prepared compounds with known δ³⁴S values.[13] A commonly-used IAEA reference material is IAEA-S-1, a silver sulfide reference material with a δ³⁴S value of -0.30‰ VCDT.[4][14] ³³S and ³⁶S abundance can also be measured relative to IAEA reference materials and reported on the VCDT scale.[13] For these isotopes, too, VCDT is established as having δ³³S and δ³⁶S values of 0‰.[13] The ³³R value of VCDT is 0.007877 and the ³⁶R value is 0.0002.[13] 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.[13]

An illustration of some common processes in the biogeochemical sulfur cycle.

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. Both sulfate and sulfide can be used by specialized microbes to obtain energy or to grow.[23] Gases including sulfur dioxide and carbonyl sulfide make up the atmospheric component of the sulfur cycle. Any process that transports or chemically transforms sulfur between these many natural materials also has the potential to fractionate sulfur isotopes.

Natural range of sulfur isotopic composition on Earth, modified and simplified from Meija et al. (2013).

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.[24] The ranges of δ³⁴S values vary across 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.[24] The range of sulfur isotope abundances in different natural materials results from the isotope fractionation associated with natural processes like the formation and modification of those materials, discussed in the next section.

Numerous natural processes are capable of fractionating sulfur isotopes. Microbes are capable of a wide variety of sulfur metabolisms, including the oxidation, reduction, and disproportionation (or simultaneous oxidation and reduction) of sulfur compounds.[1] The effect of these metabolisms on sulfur isotopic composition of the reactants and products is also highly variable, depending on the rate of relevant reactions, availability of nutrients, and other biological and environmental parameters.[25][26] 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.[25][27]

Many abiotic processes also fractionate sulfur isotopes. 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.[28] Some sulfide minerals, including pyrite and galena, can form through thermochemical sulfate reduction, a process in which seawater sulfate trapped in seafloor rock is reduced to sulfide by geological heat as the rock is buried; this process generally fractionates sulfur more strongly than gypsum formation.[29]

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.[9] Mass-anomalous fractionations are rare, but they can occur through certain photochemical reactions of gases in the atmosphere.[30][31] Studies have shown that photochemical reactions of atmospheric sulfur dioxide can cause substantial mass-anomalous fractionation of sulfur isotopes.[30][31]

Signatures of mass-anomalous sulfur isotope fractionation preserved in the rock record have been an important piece of evidence for understanding the Great Oxidation Event, the sudden rise of oxygen on the ancient Earth.[9][56] 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.[9] Multiple mechanisms have been proposed for how oxygen prevents the fingerprints of mass-anomalous fractionation from being created and preserved; nevertheless, all studies of Δ³³S and Δ³⁶S records conclude that oxygen was essentially absent from Earth's atmosphere prior to 2.45 billion years ago.[9][10][30][57][58]

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.[1] By measuring the sulfur isotopic composition of these preserved materials, scientists can reconstruct ancient biological processes and the environments where they occurred.[1] δ³⁴S values in the geologic record have been inferred to reveal the history of microbial sulfate reduction[59][60] and sulfide oxidation.[61] Paired δ³⁴S and Δ³³S records have also been used to show ancient microbial sulfur disproportionation.[62][26]

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.

Microbial dissimilatory sulfate reduction (MSR), an energy-yielding metabolism performed by bacteria in anoxic environments, is associated with an especially large fractionation factor.[1] The observed ³⁴ε_MSR values range from 0 to -65.6‰.[25][27][36][37][38][39][40][41] Many factors influence the size of this fractionation, including sulfate reduction rate,[32][37] sulfate concentration and transport,[25][41] availability of electron donors and other nutrients,[27][39][40] and physiological differences like protein expression.[42] Sulfide produced through MSR may then go on to form the mineral pyrite, preserving the ³⁴S-depleted fingerprint of MSR in sedimentary rocks.[1][54] Many studies have investigated the δ³⁴S values of ancient pyrite in order to understand past biological and environmental conditions.[1] For example, pyrite δ³⁴S records have been used to reconstruct shifts in primary productivity levels,[63] changing ocean oxygen content,[64][65] and glacial-interglacial changes in sea level and weathering.[66] Some studies compare sulfur isotopes in pyrite to a second sulfur-containing material, like sulfate or preserved organic matter.[63][64] Comparing pyrite to another material gives a fuller picture of how sulfur moved through ancient environments: it provides clues about the size of ancient ³⁴ε_MSR values and the environmental conditions controlling MSR fractionation of sulfur isotopes.[63][64]

δ³⁴S records have been used to infer changes in seawater sulfate concentrations.[67] 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.[68] δ³⁴S values of pyrite have also been applied to reconstruct the concentration of seawater sulfate, based on expected biological fractionations at low sulfate concentrations.[69][70] Both of these methods rely on assumptions about the depositional environment or the biological community, creating some uncertainty in the resulting reconstructions.[25][68]