Reef aquarists fear a number of toxic chemicals in their aquaria. Many of these, such as copper and detergents, can be avoided by controlling what is added to the aquarium. Some, however, can be generated inside the aquarium, and consequently must be controlled in other ways. One of these is hydrogen sulfide. It can be produced in an aquarium's anoxic regions, such as inside or under live rock, inside sand beds and in denitrators. At sufficiently high concentration it is not only foul smelling, but can be lethal to many marine organisms.

Being aware of when and how hydrogen sulfide can form may help aquarists avoid elevated concentrations in their aquaria, and allow them to understand how to deal with it if such events occur. This article describes in detail what hydrogen sulfide is, how it is toxic, how it forms and degrades, and suggests ways to avoid toxicity problems.

The sections are:

• The Sulfur Cycle for Reef Aquarists
• The Nature of Hydrogen Sulfide
• The Odor of Hydrogen Sulfide
• Stability of Hydrogen Sulfide in Water: Abiotic Oxidation
• Stability of Hydrogen Sulfide in Water: Iron Hydroxide Oxidation
• Stability of Hydrogen Sulfide in Water: Metal Sulfide Precipitation
• Biological Oxidation of Sulfide in Seawater
• Hydrogen Sulfide Production
• Hydrogen Sulfide in Marine Sediments
• Toxicity of Hydrogen Sulfide
• Organisms with Special Tolerance to Hydrogen Sulfide
• Hydrogen Sulfide in My Reef Aquarium
• Hydrogen Sulfide in Other Reef Aquaria
• Preventing and Dealing with Hydrogen Sulfide
• Summary
• References

The Sulfur Cycle for Reef Aquarists

Sulfur can take many forms in the marine environment. Figure 1 shows a greatly abbreviated sulfur cycle that pertains especially to processes that take place in reef aquaria. Subsequent sections of this article will expand on several of these processes, but it is worthwhile to have the whole cycle in place as a framework before such discussions.

The tour starts with sulfate (Figure 1, middle top; SO₄^--). Sulfate consists of one sulfur atom surrounded by four oxygen atoms, and it carries a charge of -2. It is the most oxidized form of sulfur in seawater and is also one of the most abundant ions, having a natural seawater concentration of about 2700 ppm. Various processes in reef aquaria can raise or lower the sulfate concentration considerably. Adding Epsom salts (magnesium sulfate heptaydrate) to boost magnesium, for example, raises sulfate more than it does magnesium. Likewise, adding calcium chloride to boost calcium reduces sulfate if salinity is maintained. Nevertheless, so much sulfate is present in seawater that such changes occur slowly over time.

Under anoxic conditions (ORP less than about 0 mV) some bacteria can use sulfate to metabolize organic material, producing hydrogen sulfide (H₂S) as a byproduct (Figure 1, middle bottom). Hydrogen sulfide is the most reduced form of sulfur present in seawater, and areas of negative redox values are known as reducing zones. In reef aquaria the reduction of sulfate can occur in many places, such as under sand or inside live rock. It also can take place in denitrators if their flow is low enough that the nitrate is depleted while organic material is still available for decomposition.

Hydrogen sulfide produced in anoxic regions can produce several possible outcomes. One is to diffuse into more oxic (aerobic) regions, even into the water column itself, where it can be toxic to organisms. In the water column it can be oxidized to sulfate or other sulfur species (elemental sulfur (S), sulfite (SO₃^-) and others; Figure 1, right). Such oxidation reactions are catalyzed both by soluble metals such as iron and by light. Hydrogen sulfide also can combine with metals such as iron (Fe⁺⁺) to precipitate as black iron sulfide (Figure 1 bottom; FeS and FeS₂). This blackness is the telltale sign of hydrogen sulfide formation that can be seen in anoxic seawater sediments, although similar appearing black precipitates may be formed from other materials.

Both hydrogen sulfide and sulfate can become sources of sulfur for the biosynthesis of organic compounds (Figure 1, left) by organisms ranging from bacteria to fish. Most proteins contain sulfur, for example, since some amino acids from which they are made contain sulfur. When organic materials that contain sulfur are metabolized, both hydrogen sulfide and sulfate can result, depending on the circumstances of that metabolism. Organic materials added to an aquarium as foods are also sources of such organic sulfur compounds.

Reef aquarists using sulfur denitrators are adding elemental sulfur (S) that is not normally present in large concentrations in oxic ocean waters. This sulfur can become oxidized by oxygen (O₂), nitrate (NO₃^-) and other oxidizing species to become sulfate.

The Nature of Hydrogen Sulfide

Hydrogen sulfide is commonly known by many other names, including hydrosulfuric acid, hydrogen sulphide, sewer gas, stink damp, sulfur hydride, sulfurated hydrogen, dihydrogen monosulfide and dihydrogen sulfide.

Hydrogen sulfide, H₂S, is analogous to the water molecule, H₂O, with the sulfur atom taking the place of the oxygen atom (Figure 2). Unlike water, hydrogen sulfide is a gas at room temperature. It can be condensed into a liquid at temperatures below -60°C, and into a solid below -85°C. The reason that it is not a liquid at room temperature like water is that the sulfur atom is much poorer at hydrogen bonding, and it is the hydrogen bonds in water that hold it in liquid form. For those chemists who use hydrogen sulfide in the laboratory, it can be purchased as a compressed gas in cylinders. Many of its physical properties can be found in this linked article.

Hydrogen sulfide has many natural and man-made sources. Volcanoes, undersea vents, swamps and other stagnant bodies of water and sulfur springs are common sources. It is also a common cause of bad breath, as it can be produced by bacteria in the human mouth at concentrations ranging up to 100 ppb. It can also comprise up to 10% of intestinal gases, and flatus can contain up to at least 18,000 ppb hydrogen sulfide. The concentration of hydrogen sulfide in typical ambient air over land is on the order of 0.1 to 0.3 ppb. In anoxic basins in the ocean, such as the Cariaco Trench or the Black Sea, hydrogen sulfide typically ranges from 0 to 200,000 ppb and is often in the 500 to 10,000 ppb range.

Hydrogen sulfide is fairly soluble in water. At 1 atmosphere H₂S partial pressure, one gram dissolves in 242 mL of freshwater at 20°C, forming a solution that has a concentration of about 4,000 ppm hydrogen sulfide. This solution is slightly acidic (pH about 4.5) because, like the water molecule, hydrogen sulfide can ionize to release H⁺:

H₂O  ßà  H⁺  +  OH^-

H₂S ßà  H⁺  +  SH^- 

Hydrogen sulfide is much more prone than water to this ionization, because the larger sulfur ion can spread out the charge to a greater extent and hence is more stable than an oxygen ion. Pure water with a pH of 7 contains equal amounts of H₂S and HS^- (hydrosulfide ion), while in order for H₂O to have equal concentrations of H₂O and OH^-, the pH needs to be above 14. At pH values above 7, HS^- dominates. At very high pH (above 11), the HS^- may ionize again to form a sulfide ion (S^--).

HS^- ßà  H⁺ +  S^--

This ionization is actually somewhat controversial. Literary sources give values for its pKa between 12 and 14, but more recent data suggests that it doesn't ionize until the effective pH is higher than 14 with lower values previously reported being erroneous due to oxidation during the experiment.

The distribution of the different forms of hydrogen sulfide is shown in Figure 3 (above) as a function of pH in freshwater (assuming the second pKa is about 12). Table 1 shows the relative ratio at pH 8.2 in freshwater. I could not find exact values for seawater, but ions usually form slightly more easily in seawater, so its proportion of H₂S at any given pH is likely lower. These forms are in rapid equilibrium, so any given S atom will convert between all of these forms many, many times each second.

Clearly it is the HS^- form that predominates in reef aquaria, but the other forms are very important. Only the H₂S form volatilizes and is detected as a foul odor. It is also likely the neutral H₂S form crosses cell membranes and enters organisms to cause potential toxicity (as is the case for ammonia, with NH₃ v. NH₄⁺, where it is the neutral NH₃ form that crosses membranes and causes toxicity).

On the other hand, even though the S^-- form is only a small fraction of the total, it is also very important as it is this form that precipitates with metals to form the black deposits characteristic of hydrogen sulfide formation in sediments. Even though it is only a small fraction of the total at any given instant, all of the sulfide can be removed from seawater by precipitation as metal sulfides if sufficient metals are available. An analogy is that only a small fraction of the water in a reef tank is inside a pump at any given instant, but over time, all of the water can pass through it and be sent to a different part of the system.

I have used this process to my advantage in some lab experiments I did years ago. After passing hydrogen sulfide gas from a cylinder through a chemical reaction, I sent the excess bubbles into a very high pH solution (> pH 14), trapping the sulfide in the water as S^--. This solution could be disposed of by a waste disposal company much more easily than if I had somehow tried to collect the gas itself.

The Odor of Hydrogen Sulfide

Hydrogen sulfide has a strong "rotten egg" smell. The odor of H₂S can be detected in the air by humans at levels as low as 0.5 to 300 ppb. The large variation in range indicates that some individuals are very much more sensitive to it than others. Interestingly, humans may become insensitive to the odor at concentrations above 100,000 ppb. For this reason, individuals working with the gas need to be aware that they may no longer smell hydrogen sulfide when it is present at life threatening concentrations. When I have used hydrogen sulfide in the laboratory, I have worn sensor badges that indicate exposure to concentrations that may be too high to smell, warning that action needs to be taken immediately (fortunately, I was very careful to work in a chemical fume hood and never smelled, nor was I exposed to any hydrogen sulfide).

When dissolved in water, the smell depends strongly on pH (which determines how much is in the volatile, hence "smellable," H₂S form). Humans often can just detect hydrogen sulfide odors when the concentration is above about 0.029 ppb in freshwater. In seawater at pH 8.2, where only 6% of the sulfide present is in H₂S, this odor threshold is likely higher, perhaps on the order of 20-fold higher (0.6 ppb). Fortunately, that threshold is below the lethal limit of many aquatic organisms (usually above 5 ppb; sometimes as high as 50,000 ppb), so odor often can be detected by humans before hydrogen sulfide rises to acute, lethal concentrations in reef aquaria.

Stability of Hydrogen Sulfide in Water: Abiotic Oxidation

In its gas phase, hydrogen sulfide is unstable toward oxygen and consequently, hydrogen sulfide is a flammable gas. While hydrogen sulfide in water is stable under anoxic and metal free conditions, in the presence of oxygen (O₂) it is unstable. When bubbled into freshwater in the presence of oxygen, hydrogen sulfide quickly reacts with oxygen resulting in the precipitation of elemental sulfur:

2 H₂S  + O₂  à  2 H₂O + 2 S

Under conditions typically found in seawater, hydrogen sulfide reacts with oxygen to form a variety of oxidized sulfur species (sulfate, SO₄^--; sulfite, SO₃^-- and thiosulfate, ^-SSO₃^-, which looks like sulfate except that one oxygen atom is replaced by an additional sulfur atom).¹ The percentage of each of these species that is formed depends on the salinity, the pH and the temperature, but the predominant product is typically sulfate.

H₂S  +  O₂  à   [SO₄^--; SO₃^--; ^-SSO₃^-]

The rate of the oxidation of sulfide in this fashion depends linearly on the concentration of both sulfide and oxygen, and can be catalyzed by certain metals, such as iron.¹ Typically the half life of hydrogen sulfide under such conditions is less than a day, and can be only a couple of hours. In a reef aquarium where iron is dosed, the half life may be much lower as the added iron may accelerate the oxidation to even higher rates, potentially protecting organisms.

Hydrogen sulfide in seawater is also oxidized photochemically. When 300 ppb hydrogen sulfide was added to Biscayne Bay and Gulf Stream waters off of Florida, the half life of the sulfide was 49 and 147 minutes, respectively. Such studies found that sunlight (both ultraviolet and visible) was able to significantly accelerate the oxidation.²

Stability of Hydrogen Sulfide in Water: Iron Hydroxide Oxidation

It has been shown that particles of iron oxide/hydroxide can react with hydrogen sulfide fairly rapidly to produce elemental sulfur and reduced metals. For example, Fe⁺⁺⁺ (ferric iron) is reduced to Fe⁺⁺ (ferrous iron) as the sulfide is oxidized to elemental sulfur. The process involves sulfide coming to the surface and giving up electrons to the metal ions that take them up.³ The rate of this process is decreased by sulfate, and less so by magnesium and calcium ions. It is strongly decreased by phosphate and silicate that bind to the particles' surfaces. Certain organic chelators (EDTA and TRIS, for example) accelerate this process by speeding the cycling between Fe⁺⁺ and Fe⁺⁺⁺.⁴ It has been claimed that this process is rapid enough to purify seawater polluted with hydrogen sulfide, as in mariculture systems.⁵

Since many reef aquarists have sufficient amounts of iron oxide/hydroxides in their systems to bind iron, this may be an important mechanism for hydrogen sulfide detoxification in some reef aquaria.

Stability of Hydrogen Sulfide in Water: Metal Sulfide Precipitation

Sulfide in seawater is also unstable toward precipitation with certain metals. The black deposits often seen in anoxic sediments are typically metal sulfides, especially ferrous sulfide (FeS) and pyrite (FeS₂), with much smaller amounts of copper, manganese, zinc, nickel and cobalt sulfides. The exact processes whereby these metal sulfides form in marine sediments and elsewhere is complicated and still under study.⁶ In some areas, like the Orca basin in the Gulf of Mexico, deposited iron sulfides make up as much as 0.7% of the sediments' mass.⁷ So, iron sulfides are not necessarily a trace component.

Biological Oxidation of Sulfide in Seawater

Hydrogen sulfide can also be taken up by bacteria, and oxidized under aerobic conditions back to sulfate. In this process the bacteria gain energy, much as other organisms gain energy by oxidizing organic (carbon) compounds. Bacteria can also use manganese oxide instead of O₂ to produce sulfate from hydrogen sulfide, so the process can occur even in anoxic conditions. The process does not occur in a single chemical step, but usually involves such intermediates as thiosulfate (^-SSO₃^-).

Hydrogen Sulfide Production

All organisms oxidize organic compounds in various ways to gain energy or to produce new organic biomolecules that they require. Most of the large organisms kept in reef aquaria (fish, corals, algae, etc.) perform this oxidation using oxygen (O₂) as the ultimate electron acceptor in the process. For example, the oxidation of methane, CH₄, to CO₂, can be done with oxygen:

CH₄  +  2O₂  à  CO₂  + 2H₂O

The reason that oxygen is called an electron acceptor is that in the reaction above, electrons are transferred from the carbon and hydrogen atoms to the oxygen atoms.

Organisms typically carry out more complicated oxidations, but the process is the same. The equation below shows the overall chemical reaction involved in the oxidation of organic molecules with oxygen, shown specifically for an organic molecule representing typical plankton.⁸

(CH₂O)₁₀₆(NH₃)₁₆(H₃PO₄) + 138 O₂ à 106 CO₂ + 122 H₂O + 19 H⁺ + PO₄^--- + 16 NO₃^-

Under conditions where O₂ is limited in supply, organisms need to turn to other electron acceptors. These include nitrate (the way that nitrate is reduced by a deep sand bed), various metals such as iron and manganese, and most important to the context of this article, sulfate. For organisms that use sulfate instead of oxygen, the oxidation can be described as:

CH₄  +  SO₄^--  à  CO₂  +  S^--  +  2 H₂O   (which is the same as HCO₃^- +  HS^- + H₂O)

The more generalized reaction for a typical "organic" can be described as:

(CH₂O)₁₀₆(NH₃)₁₆(H₃PO₄) + 53 SO₄^-- à 56 CO₂ + 50 HCO₃^- + 53 HS^- + 16 NH₃ + 53 H₂O + PO₄^---

Note that the ammonia produced is not oxidized to nitrate under these conditions as it is under aerobic conditions. Each of these processes produces different amounts of energy for the organisms involved. Oxidation with O₂ produces the most energy, followed by nitrate, then manganese, iron, sulfate and finally carbon dioxide with methane as the product. Organisms (or more correctly, ecosystems that evolve containing many different organisms) often get as much energy as they can from a food supply, so until the O₂ runs out, few organisms carry out anything other than oxidation processes. The other processes are each carried out in turn, either chronologically in certain situations, or with depth into a substrate. Sulfate is one of the last usable electron acceptors available in seawater. However, it is also available in far higher concentration (2700 ppm) than any of the other acceptors (which are often sub ppm), so it can sometimes be used in parallel with all the other electron acceptors besides O₂.

These processes are typically carried out by bacteria and archaea under low oxygen conditions. They take up sulfate via transporters in their cell membranes, and then use it internally in a series of separate chemical processes ending with it being transformed into hydrogen sulfide.

In addition to these biological processes, purely chemical reactions also produce hydrogen sulfide in the ocean. The heat of hydrothermal vents, for example, can drive the reaction between organic materials and sulfate to produce hydrogen sulfide. While the process can theoretically proceed at as little as 25°C, it is so slow at that temperature as to be unimportant. At 200°C, typical reactions along these lines can take months to years, and it has not been demonstrated at less than 125°C. Such reactions should consequently have minimal importance in most reef aquaria.

Hydrogen Sulfide in Marine Sediments

The study of the chemistry or the pore water in marine sediments is widespread, extensive and ongoing. It is beyond the scope of this article to cover this topic, but a few comments about hydrogen sulfide in this environment are warranted as sediments are the most likely place for hydrogen sulfide to be produced in a reef aquarium.

Because of the processes described above, marine sediments often accumulate hydrogen sulfide and deplete in sulfate. This zone often starts a few centimeters below the surface, and can extend up to a meter or more before the sulfate is fully depleted. Below that depth, other processes take place, such as methane production. These additional processes are also beyond the scope of this article.

In general, as seawater becomes depleted of oxygen, a series of chemical transformations takes place, largely due to biological activity continuing to consume oxygen and other electron acceptors. There is a specific order of usage of electron acceptors that can result in a layering of chemical processes with depth, either into sediments, or sometimes into anoxic waters in enclosed basins.

As mentioned above, this order of electron acceptors used to oxidize organic material is oxygen (O₂), then nitrate (NO₃^-), then manganese (Mn⁺⁺⁺⁺), then iron (Fe⁺⁺), then sulfate (SO₄^--). Researchers can plot the concentrations of these chemicals as a function of depth, and can also associate an ORP with each transition, although some overlap of the chemistries takes place in each zone. The oxygen zone has an ORP of 0 to 600 mV, the nitrate zone is -150 to 550 mV, the manganese zone is -50 to 400 mV, the iron zone is -700 to -150 mV and the sulfate reduction zone is -850 to 0 mV. Consequently, if hydrogen sulfide is being formed from sulfate, the ORP is likely below 0 mv in that region.

Exactly what processes take place at what depths in sediments depends on many factors, such as the nature of the sediments themselves (size distribution and chemical makeup), the amount of organic material being deposited and the temperature. In one study of the sediments below seawater fish farms, the hydrogen sulfide level in the sediments appears to vary with the season, and peaks as high as 70,000 ppb in the Fall.⁹ Another study showed that even though sulfate reduction took place maximally at 1-2 cm depth, free hydrogen sulfide was present only below a depth of 6-7 cm, above which Fe⁺⁺ was available for precipitation as iron sulfides.¹⁰

Toxicity of Hydrogen Sulfide

Hydrogen sulfide is toxic to a wide range of organisms, including people. That fact has been known for hundreds of years. It is just now becoming clear, however, that hydrogen sulfide also appears to play important roles in normal biochemical processes in animals. Neurons and muscles, for example, may use it in various ways, but exactly how this takes place has not been elucidated.¹¹

One way that hydrogen sulfide exerts its toxicity is by inhibiting a mitochondrial enzyme called cytochrome c oxidase. It can be inhibited at hydrogen sulfide levels in solution as low as 30 ppb.¹² Such inhibition limits the ability of mitochondria to produce energy for cells. Another enzyme, catalase, is inhibited at concentrations of 6,000 ppb.¹² Other mechanisms of toxicity are also likely, and have recently been studied.¹³

Unfortunately for many organisms, their internal hydrogen sulfide concentrations can be significantly higher than its external concentration. For example, the concentration of hydrogen sulfide in the hemolymph of the marine worm Halicryptus spinulosus can be three times higher than the external concentration.¹²

A review of the health effects of hydrogen sulfide in general is given in this online review:

Toxicological Profile for Hydrogen Sulfide

Tables 2 and 3 summarize some data in the scientific literature for the lethality of hydrogen sulfide and also for sodium sulfide (which dissolves to form a sodium ion and a sulfide ion) to a variety of aquatic organisms. The data are given as the LC₅₀, which is the concentration at which half of the organisms die when exposed over a period of an hour to a few days. Briefer exposures are typically less likely to cause death than are longer exposures to the same concentration. For example, the LC₅₀ for Gammarus pseudolimnaeus after 96 hours of exposure was 22 ppb, but in a longer test of 105 days, the concentration needed to be below 2 ppb to ensure survival.¹⁴ Clearly, the lethal concentrations vary tremendously from organism to organism, ranging from 7 ppb for Brown Trout to 750,000 ppb for the Western Mosquitofish.

The toxicity of hydrogen sulfide to many organisms varies with salinity, pH, temperature and other factors. The clam Meretrix lusoria, for example, is most susceptible to H₂S poisoning at 15-20 ppt salinity, pH 7.5-8.5 and temperatures above 25°C (which likely relates to the clam's changing metabolic activity as conditions vary from these).¹⁵ A similar temperature effect was observed in goldfish (Carassius auratus).¹⁶

Most of the data in Tables 2 and 3 were obtained from a free online database (PAN Pesticide Database; www.pesticideinfo.org) at the following pages:

Hydrogen Sulfide
http://www.pesticideinfo.org/Detail_Chemical.jsp?Rec_Id=PC39183#Ecotoxicity

Sodium Sulfide
http://www.pesticideinfo.org/Detail_Chemical.jsp?Rec_Id=PC38989#Ecotoxicity