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
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,
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.
Figure 1. Simplified sulfur cycle in reef aquaria.
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
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
ocean waters. This sulfur can become oxidized by oxygen (O2),
and other oxidizing species to become sulfate.
The Nature of 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
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
Figure 2. Space-filling molecular models showing the
water (H2O; left)
to the larger hydrogen sulfide (H2S;
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+ + OH-
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--).
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.
Figure 3. Speciation of hydrogen sulfide as a function
of pH in freshwater.
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.
1. Speciation of hydrogen sulfide in fresh water at
Clearly it is the HS- form
that predominates in reef aquaria, but the other forms are
very important. Only the H2S
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
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:
+ 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
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.
+ O2 à
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++
It has been claimed that this process is rapid enough to purify
seawater polluted with hydrogen sulfide, as in mariculture
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
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
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
+ 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
+ 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:
+ 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:
+ 53 SO4-- à
56 CO2 + 50 HCO3- + 53
HS- + 16 NH3 + 53 H2O +
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
These processes are typically carried out by bacteria and
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
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
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
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:
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
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:
2. Toxicity of hydrogen sulfide to various aquatic
Oriental river shrimp
Rainbow trout, Donaldson trout
Stizostedion vitreum vitreum
Scud = freshwater amphipod
Oriental river shrimp
3. Toxicity of sodium sulfide (Na2S)
to various aquatic organisms.
Dungeness or edible crab
Giant river prawn
White cloud mountain minnow
Common bay mussel, blue mussel
Organisms with Special Tolerance
to Hydrogen Sulfide
Many marine species live in close
proximity to sediments that often contain hydrogen sulfide.
Some even live in them. From the data in Tables 2
and 3, it is clear that the range of
susceptibility to hydrogen sulfide poisoning is huge, and
those species more prone to natural exposure often have higher
tolerance. How do they accomplish that? Likely by more ways
than are presently known, but even so, some of these organisms'
mechanisms are known.
In terrestrial mammals (dogs, cats, rabbits), hydrogen sulfide
is primarily detoxified by reaction with oxyhemoglobin to
form colloidal sulfur, hemoglobin and water.17
However, if hydrogen sulfide concentrations are high, or if
oxygen concentrations are low, that mechanism is not adequate.
In fact, conditions of anoxia increase the negative effects
of hydrogen sulfide on many aquatic organisms. 15,16,18
In the marine worm Urechis
caupo, for example, it has been shown that excess
hydrogen sulfide is oxidized to thiosulfate (-SSO3-).19
The thiosulfate is then allowed to passively diffuse out of
the organism's hindgut. Interestingly, this worm also appears
to make hydrogen sulfide out of sulfur-containing amino acids
and use it as a gasotransmitter, controlling its body wall's
muscle tone.20 Whether this
mechanism also helps it limit exposure to ambient hydrogen
sulfide is unclear.
The marine worm Halicryptus spinulosus has an even
more elaborate system. During exposure to hydrogen sulfide
under oxic conditions, it oxidizes the hydrogen sulfide to
thiosulfate, just as the Urechis caupo does. Under
anoxic conditions, it has a multi-pronged strategy. Its first
defense is to supply iron for precipitation of iron sulfide
on its surface and in its blood. Under these conditions, the
animal's surface and its blood blacken considerably, but this
process is reversible once oxic conditions return. Finally,
these worms apparently bind sulfide to an unidentified chemical
in the hemolymph, providing additional protection for its
Another way that organisms detoxify hydrogen sulfide is to
incorporate bacterial symbionts that themselves detoxify the
sulfide. Some bivalves and annelids, for example, have special
hemoglobin structures that bind sulfide, thereby detoxifying
it, and that also serve to transport it to such bacteria that
can remove the sulfide from the hemoglobin.21
Hydrogen Sulfide in My Reef Aquarium
A number of aquarists have reported
black metal sulfide deposits in their reef aquaria, especially
in sand beds. I recently took down a reef aquarium that had
been in operation for 10 years. It was doing just fine (Figure
4 below) and had never suffered any decline that some aquarists
would call "old tank syndrome." I wanted to replace
the stand, however, with a more attractive custom stand, and
to do so required taking it all apart. So, I took the opportunity
to replace the glass aquarium as well.
Figure 4. My ten-year-old reef aquarium that I took
down in October, 2005 to replace the stand.
The aquarium had a bed of 1-2 inches of fine oolitic aragonite
sand and contained wild Florida live rock (not aquacultured)
placed in the tank when it was started. As I dug through the
sand, I found no black areas, and detected no unpleasant odors.
None of the rock had any apparent grey or black discoloration.
All rocks were a uniform tan color on areas not exposed to
light. The sand was put into a bucket and remained in the
garage for two weeks. When I then dumped the sand out behind
the garage, it was grey and it stank to high heaven. Clearly
the processes leading to hydrogen sulfide formation were not
taking place in the sand bed in the tank, but that same sand
quickly became anoxic when removed from the aquarium.
As an experiment, I also took some samples from one of my
refugia. This refugium had about six inches of fine oolitic
aragonite sand on its bottom, and has live rock and macroalgae
above that. It has been in place for several years. I took
samples by pressing glass test tubes into the sand, removing
cores of sand that were 1-4 inches long. None showed any signs
of discoloration or odor. I then set one of these cores aside.
To a second core sample I added some flake food on the top
of the core while still in the tube, then added another core
on top of it in the same test tube. This tube was also set
After two weeks I examined them. I expected more discoloration
in the tube containing the flake food, as it provides a lot
of organic material to decompose, but that really didn't seem
to be the case, at least over the two weeks involved. There
was discoloration around some of the flakes, but mostly the
flakes seem unchanged with just a few of them being discolored
and with other scattered black spots throughout the sample
(Figure 5). Even the unaltered core smelled bad and had black
discoloration in distinct areas (Figure 6). Apparently, plenty
of organic material is already in these sands to drive anoxia
IF they are removed from the system so that oxygen
cannot get to them.
Figures 5 & 6. Sand that had been in my aquarium
for many years, collected by coring a deep sand bed.
Flake food was added on top of the sand sample after
coring, and then another inch of live sand was added
on top of that. It was set aside for two weeks before
the picture was taken.
I had also heard that live rock sometimes can be shown to
have internal sulfide deposits. To test this idea, I decided
to break open some of the live rock that had been in the tank.
Figure 7 shows a typical rock that I examined. As mentioned
above, it was wild rock collected from Florida about 10 years
ago. It was a uniform tan color on the outside, and had not
been buried under the sand. Upon breaking it open with a hammer
(Figure 8), it is clear that there are some grayish sulfide
deposits in it (Figure 9), but they
do not occupy the rock's entire interior. These deposits did
not smell, but are probably metal sulfides nevertheless.
Figures 7 & 8. Left: A live rock collected
wild in Florida more than ten years ago that had been in my
aquarium since collection. Right: The same rock as
in Figure 7, now broken open.
Figure 9. A close-up of one of the fractured rocks
Figure 8. The dark band suggests metal sulfide deposits.
Hydrogen Sulfide in Other Reef
Other aquarists have not been as fortunate
as I have with respect to hydrogen sulfide formation, especially
in their sand beds. Some aquarists have found anoxic regions
in deep sand beds that reek when disturbed. Sometimes aquarists
have run into problems when trying to remove a sand bed containing
hydrogen sulfide. In some cases, corals and other organisms
have died as a result. Whether these resulting deaths were
from hydrogen sulfide or other causes is impossible to say,
but it is a reasonable hypothesis, and one that can provide
some potential preventive measures that will be discussed
It is not, however, always trivial to distinguish black sulfide
deposits from very dark forms of algae or cyanobacteria, or
simply old organic debris. In totally dark regions (such as
several inches below a sand bed), black deposits are very
likely due to metal sulfide formation. Modern furniture, new
prices, and lower prices in Early Settler Catalogue are here. But closer to the surface,
such as along the glass under the sand's surface, it often
can be difficult to identify by looking at a picture alone.
In soliciting pictures for this article, a great many of those
submitted looked more like algae to me than metal sulfides.
Figure 10 shows a black patch in a marine aquarium that may
likely be sulfide deposits.
Figure 10. A dark deposit in the sand bed of another
that may be metal sulfide.
Preventing and Dealing with Hydrogen
Aquarists sometimes have to deal with
situations where hydrogen sulfide exposure is a possibility.
The suggestions below may be useful in preventing or correcting
1. Avoid burying organic materials under sand or rocks. This
organic material decays, and that decay can drive the formation
of hydrogen sulfide. Burying live rock under sand, for example,
will likely kill many organisms and when they decay, H2S
may be formed.
2. When using a denitrator, be careful not to set its flow
too slow so that nitrate is eliminated and sulfate reduction
takes over. The ORP
in the denitrator may or may not be a useful guide here, depending
on the setup.
3. Do not store live sand for extended periods without placing
it in circulating aerated water.
4. Do not add organisms (such as certain fish) that dig up
sand beds if there is significant potential for hydrogen sulfide
5. Be careful that equipment failures (such as a powerhead
falling off the side of an aquarium onto the sand) will not
disturb sand beds if sulfide deposits are suspected.
6. Do not stop the flow of aquarium water inside piping or
other closed systems for more than an hour or two. If it is
necessary to stop the flow for a longer time, collect the
water rather than sending it back into the tank.
7. Do not store tank water or natural seawater without stirring
and aeration for more than a few hours. Raw artificial seawater
made with pure fresh water does not have this concern, as
there are no organics in it to decay. It can be stored unstirred
for as long as desired.
8. If an anoxic sand bed needs to be removed from a reef
aquarium, and there are organisms that cannot be relocated
out of harms way, the following precautions may be useful
based on the principles detailed in previous sections, although
I've not tested any to see how effective they are:
A. Remove delicate organisms from the tank system, if possible.
B. Perform the change when the lights are as bright as possible,
preferably near the end of the light cycle. The lights drive
the O2 concentration higher,
speeding the oxidative removal of hydrogen sulfide, and
the light itself will catalyze the oxidation of H2S.
C. Maximize aeration. A high oxygen level drives hydrogen
sulfide oxidation, and high aeration will drive some off
as volatile H2S gas.
D. Add an iron supplement to help catalyze oxidation of
hydrogen sulfide and the precipitation of ferrous and/or
ferric sulfide. Use one chelated to an organic; either ferrous
or ferric iron will work.
E. Pass the water over iron oxide/hydroxide (GFO) to convert
hydrogen sulfide to elemental sulfur.
F. Pass the water over activated carbon, which may bind
some sulfide, and may also catalyze the oxidation. If forced
to choose between carbon and GFO, I'd pick the GFO media.
Hydrogen sulfide is something that
most reef aquarists eventually encounter. Some unfortunate
aquarists encounter it on the first day of their exposure
to reef aquaria when their live rock arrives and dead organisms
are decaying. From that day on, aquarists should be aware
of the risk of generating hydrogen sulfide, especially in
Special thanks to Jens Kallmeyer for supplying a copy of
his doctoral thesis "Sulfate Reduction in the Deep Biosphere,"
which provided an extensive background on a variety of the
processes described in this article.