Sulfur Cycle

For a high-level tour of how every element moves through the tank, start with Nutrient Cycling. Sulfur is the element behind the most feared failure mode in a planted tank: the black, rotten-egg-smelling substrate that quietly poisons shrimp from below.

Sulfur leads a double life

Sulfur is two elements wearing one coat. By the demands of biology it is a quiet macronutrient — every protein needs the sulfur-bearing amino acids cysteine and methionine, every chloroplast runs on iron-sulfur clusters, and the cell's main antioxidant is built around a sulfur atom. Cells need it in roughly the same proportion as potassium, far more than any trace metal. In a normal tank that demand is trivially met: tap water already carries plenty of sulfate, and producers simply take what they need.

By chemistry, though, sulfur behaves like iron. It is redox-active, it cycles through the oxygen-free sediment, and — uniquely among the major nutrients except nitrogen — it can leave the tank entirely as a gas. This is the side that matters to aquarists, because it is the chemistry that drives old tank syndrome: the slow drift of a neglected, heavily-detritus-laden tank toward a black substrate and dissolved hydrogen sulfide that can kill sensitive invertebrates with no obvious water-test warning. Every step of that story is a sulfur step, so even though most healthy tanks never see a whiff of sulfide, the cycle earns its own page.

The two faces of dissolved sulfur

Dissolved sulfur exists in the tank in two completely different chemical states, and the model keeps them strictly separate because biology and chemistry treat them as different substances.

Sulfate is the oxidized, harmless, useful form — the form plants and microbes actually eat. Producers reduce sulfate inside their cells, step by step, into the cysteine they build into protein. Tap water typically carries somewhere between a couple and twenty-odd milligrams of sulfur per litre as sulfate, comparable to potassium and three orders of magnitude more than iron — so much that producers are essentially never sulfur-limited in a real aquarium. The model reflects this: it lets organisms draw sulfate freely without any scarcity throttle, because at realistic concentrations a throttle would never engage.

Sulfide (bisulfide, and its volatile cousin hydrogen sulfide) is the reduced, toxic form — the rotten-egg gas. It is produced by sulfate-reducing bacteria deep in anoxic sediment, and biology cannot eat it on aquarium timescales. Keeping these two apart matters: if the model let a cell build its amino acids out of sulfide it would be representing cellular biochemistry completely wrongly. So sulfide is invisible to the uptake machinery and reachable only by the chemistry that oxidizes it, traps it, or vents it.

The sediment carries the same split — pore-water sulfate that diffuses down and feeds reduction, and pore-water sulfide that precipitates out as a black mineral — plus the buried iron-sulfide buffer that is the heart of this whole story. In a sealed tank, hydrogen sulfide can also build up in the air space above the water.

Sulfate reduction: where the trouble starts

Deep in the substrate, where oxygen has run out and the easier electron acceptors (oxygen, nitrate, iron oxides) have been used up, sulfate-reducing bacteria turn to sulfate to keep burning organic carbon. They strip oxygen atoms off the sulfate and leave sulfide behind:

sulfate + organic carbon → sulfide + dissolved inorganic carbon

This only runs in genuine anoxia and only where there is enough buried detritus to fuel it, so the model gates it sharply on oxygen-free conditions and on settled-detritus carbon — exactly the conditions that develop in an aging, organic-rich substrate. The pace is slow: at typical hard-water sulfate levels and full anoxia it amounts to a few milligrams of sulfur reduced per litre of pore water per month (Canfield 1991). That slowness is the point — it is why the classic sulfide event takes the better part of a year to develop in a Walstad-style tank rather than days. The exact rate constants and gating thresholds are tabulated in the Parameter Reference.

Why healthy tanks never smell of sulfide

A tank can be quietly reducing sulfate in its substrate for months and still show zero sulfide in the water — because sulfide has to survive two gauntlets before it can reach the open water, and in a healthy tank it survives neither.

The first gauntlet is oxygen. Any sulfide that diffuses up into oxygenated water is re-oxidized back to sulfate within minutes (Millero 1991). The model runs this reaction fast, as it genuinely is, so water-column sulfide is held at essentially zero in any tank with a normally oxygenated water column. The same thing happens at the thin oxygenated skin at the top of the substrate — a "cryptic" oxidation zone that re-oxidizes pore sulfide before it can escape. This is why planted tanks whose rooted plants leak a little oxygen from their roots almost never develop sulfide problems, even when their substrates are reducing sulfate vigorously underneath.

The second gauntlet — the decisive one — is iron.

Iron: the substrate's sulfide buffer

When dissolved iron in the pore water meets sulfide, the two snap together almost instantly into black iron sulfide (mackinawite) and drop out of solution:

iron + sulfide → iron sulfide (black solid)

This single reaction is what protects a tank. As long as the substrate still has reducible iron in it, every molecule of sulfide is captured the moment it forms and locked away as a harmless black mineral — the same blackening you see when you dig into an old substrate. The iron stock in the soil therefore acts as a sulfide buffer that can soak up months, sometimes years, of sulfate reduction before it is exhausted. An iron-rich aquasoil has a deep buffer; a lean, sandy, or peaty substrate has a shallow one.

When that buffer finally saturates — when there is no iron left to bind the incoming sulfide — the character of the substrate changes. Pore sulfide starts to climb, then to diffuse upward faster than the water column can oxidize it, and free hydrogen sulfide begins to accumulate and escape. This is old tank syndrome: the rotten-egg smell and the unexplained livestock losses of a substrate that has used up its iron defence. The buffer can be slowly recharged — re-oxidizing the iron sulfide back to iron and sulfate when oxygen returns to the substrate, whether from a deep gravel vacuum or from oxygen leaking off plant roots — but this is a days-to-weeks process (Berner 1981), so a single substrate disturbance does not instantly reverse a long sulfidic episode. The same coupling, told from the iron side, is in the Iron Cycle.

Why pH decides how toxic sulfide is

Not all dissolved sulfide is dangerous. Only the un-ionized form — molecular hydrogen sulfide, the part that smells and the part that crosses biological membranes — is toxic and volatile. How much of the dissolved sulfide is in that form depends almost entirely on pH, because sulfide is a weak acid.

In acidic water most of the sulfide is the toxic, volatile form; around neutral pH it is split roughly half and half; in alkaline water only a small fraction is toxic. At pH 6 the large majority of dissolved sulfide is the dangerous gas; by pH 8 only about a tenth of it is. This is why a sulfide event in a soft, acidic tank is so much more dangerous than the same amount of sulfide in a hard, alkaline one — and it mirrors exactly the way pH governs ammonia toxicity, just in the opposite direction.

In an open-topped tank, the hydrogen sulfide that does form escapes to the room air (the smell you notice) and the dissolved level stays low. In a sealed jar it instead builds up in the headspace and equilibrates with the water, so a sealed sulfidic system can develop a genuinely toxic atmosphere above the water line.

Who dies first

As a sulfide event develops, the animals do not all succumb at once — they drop out in a characteristic order set by their sensitivity. Shrimp are the most vulnerable: dwarf shrimp like Neocaridina start dying at sulfide concentrations of only around ten micrograms per litre. Daphnia and copepods follow at somewhat higher levels, ostracods are more tolerant still, and snails are the hardiest, persisting after everything else has gone. This staggered die-off — shrimp first, snails last — is one of the clearest fingerprints of a sulfide problem, and it is why "the shrimp died but the snails are fine" is such a telling symptom. The per-species toxicity thresholds are listed in the Parameter Reference.

Where sulfur comes from, and dosing

Almost all of a tank's sulfur arrives dissolved in the source water, and the range across water types is wide:

Hobbyists almost never set out to dose sulfur for its own sake. Instead it rides along as the counter-ion of other supplements: Epsom salt (magnesium sulfate), sulphate-of-potash (potassium sulfate), and gypsum (calcium sulfate) all deliver sulfate alongside the nutrient you actually wanted. The model accounts for this — dosing a metal in its sulfate form, or a potassium or magnesium sulfate salt, credits the sulfate that comes with it — though at hobbyist dose rates the amount is tiny next to the tap-water baseline, and there is also a direct sulfur-dosing option for anyone who wants to add sulfate explicitly. Unlike iron and copper, sulfate needs no chelation: it is dissolved and ready to use exactly as it arrives.

How sulfur sits in the budget

Sulfur is conserved across the tank, with only a few honest exits that the mass-balance diagnostics track: sulfate (and any suspended sulfide) carried out in a water change, hydrogen sulfide vented to the room in an open-topped tank, and iron sulfide buried indefinitely in the substrate — which is not really a loss, just a slow reservoir sitting in the sediment. In a closed jar with no water changes, sulfur closes to within a fraction of a percent over a year-long run; drift larger than that points at a bookkeeping bug rather than real chemistry.

How sulfur interacts with the other cycles

• Iron — the headline coupling. Pore iron and pore sulfide precipitate together as black iron sulfide, so the substrate's iron stock is the buffer that holds off a sulfide event; when the buffer saturates, sulfide breaks through. See Iron Cycle.
• Carbon — sulfate reduction is one of the anaerobic ways buried organic carbon gets oxidized, working in parallel with denitrification and iron reduction. In an aged, organic-rich sediment it can become a real contributor to pore-water carbon dioxide.
• Oxygen — re-oxidizing sulfide (and iron sulfide) back to sulfate consumes oxygen. During an active sulfide event this becomes a genuine oxygen drain on the water column, feeding the low-oxygen conditions that go hand in hand with old tank syndrome.
• Phosphorus — an indirect link only: sulfate reduction eats settled detritus, which releases the phosphorus (and nitrogen, iron, and trace metals) bound up in that organic matter along the way.

Further reading

• Iron Cycle — the iron-sulfide buffer that holds off old tank syndrome, told from the iron side
• Phosphorus Cycle — the parallel sediment redox chemistry that releases phosphorus
• Soil and Pore Water — the sediment compartment where sulfate reduction and iron-sulfide burial happen
• Stability and Failure — where old tank syndrome sits among the ways a tank can fail
• Parameter Reference — every rate constant, half-saturation, and toxicity threshold behind this page, with citations

Key references

• Berner, R.A. (1981). Sedimentary pyrite formation: an update. Geochimica et Cosmochimica Acta 48, 605–615.
• Canfield, D.E. (1991). Sulfate reduction in deep-sea sediments. American Journal of Science 291, 177–188.
• Jørgensen, B.B. (1982). Mineralization of organic matter in the sea bed — the role of sulphate reduction. Nature 296, 643–645.
• Millero, F.J. (1991). The oxidation of H₂S in the ocean. Geochimica et Cosmochimica Acta 55, 1947–1955.
• Sander, R. (2015). Compilation of Henry's law constants for water as solvent. Atmospheric Chemistry and Physics 15, 4399–4981.
• Sterner, R.W. & Elser, J.J. (2002). Ecological Stoichiometry: The Biology of Elements from Molecules to the Biosphere. Princeton University Press.
• Stumm, W. & Morgan, J.J. (1996). Aquatic Chemistry, 3rd ed. Wiley-Interscience.