Redox Chemistry in Aquariums: Electron Acceptors, Substrate Zones and Sulfide

How oxidation-reduction works in tank water and substrate: the redox ladder from oxygen to sulfate, layered zones in sediment, hydrogen sulfide formation and what ORP values mean.

Oxidation and reduction in aquarium water

Oxidation is the loss of electrons, or an increase in oxidation state; reduction is the gain of electrons, or a decrease in oxidation state. The two always happen together: the reducing agent gives up electrons and is oxidised, while the oxidising agent accepts them and is reduced. A reduced species and its oxidised form make a redox couple, for example Fe2+/Fe3+.

In water and sediment, redox reactions are mostly driven by microorganisms. They gain energy by moving electrons from a donor, usually organic carbon, to an acceptor, usually an inorganic species. In a tank the donors are uneaten food, faeces, dead plant tissue and other organic debris. Dissolved oxygen is the first acceptor used: O2 + 4H+ + 4e- → 2H2O.

The redox ladder

Electron acceptors do not all yield the same energy. Microbes use the acceptor that gives the most energy first and move to the next one only when it runs out. The US Geological Survey (USGS) describes the usual succession as oxygen, then nitrate, manganese(IV), iron(III), sulfate and finally carbon dioxide. Wikipedia calls this ordering a redox ladder or redox gradient. It is not strictly thermodynamic: ecological and physiological factors also affect which microbes are active where.

1. Oxygen reduction (aerobic respiration): O2 + 4H+ + 4e- → 2H2O.
2. Nitrate reduction, either to nitrogen gas by denitrification (2NO3- + 12H+ + 10e- → N2 + 6H2O) or to ammonium (NO3- + 10H+ + 8e- → NH4+ + 3H2O).
3. Manganese reduction: solid MnO2 + 4H+ + 2e- → dissolved Mn2+ + 2H2O.
4. Iron reduction: solid Fe(OH)3 + H+ + e- → dissolved Fe2+ + H2O.
5. Sulfate reduction: SO4 2- + 9H+ + 8e- → HS- + 4H2O.
6. Methanogenesis: CO2 + 8H+ + 8e- → CH4 + 2H2O.

Each rung leaves a chemical trace: dissolved manganese, ferrous iron, sulfide and methane appear only after the acceptors above them are used up. For groundwater, the USGS framework calls water oxic at 0.5 mg/L dissolved oxygen or more, and reads dissolved manganese of at least 0.05 mg/L or iron of at least 0.1 mg/L as signs of manganese or iron reduction. These thresholds were set for aquifer studies, not aquaria.

Redox zones in the substrate

In a sealed container of sediment, the ladder plays out over time. In a pond or tank, oxygen keeps diffusing in from above, so the same reactions are layered by depth instead. In aquaculture ponds, pore water usually has no dissolved oxygen below a few millimetres. Denitrification sits at the top of the anaerobic layer, with iron and manganese reduction, sulfate reduction and methanogenesis below it. Organic matter breaks down fastest at the oxic-anoxic boundary.

How deep these layers sit depends on the organic load. Freshly settled, easily decomposed material such as uneaten feed, faeces and dead plankton creates most of the sediment's oxygen demand, and even a small amount can make the surface anaerobic. In nutrient-poor lakes the redox potential may not drop to 0.2 V until 5 to 10 cm down, compared with a few millimetres in productive ponds.

The layers can be seen. Aerobic sediment keeps its natural brown, yellow or grey colour. Anaerobic sediment turns dark grey or black, because ferric iron (Fe3+) is reduced to ferrous iron (Fe2+) at a redox potential of about 0.2 V.

Sulfate reduction and hydrogen sulfide

Sulfate-reducing bacteria and archaea use sulfate (SO4 2-) as their final electron acceptor and release hydrogen sulfide as waste. As electron donors they oxidise organic acids such as lactate and acetate, alcohols, or hydrogen gas, competing for hydrogen with methane-producing microbes. The rotten-egg smell of hydrogen sulfide is a sign that they are active. According to Boyd, sulfate reduction starts only once the redox potential falls to about 0.0 to 0.1 V.

Seawater contains about 2,700 mg/L of sulfate, so hydrogen sulfide problems are more likely in brackish and marine systems than in freshwater. Sulfide reacts with dissolved metals to form insoluble, often dark sulfides. In sediment, ferrous iron and sulfide combine into iron sulfide minerals such as pyrite, which stay stable as long as conditions remain anaerobic.

Hydrogen sulfide is a weak acid, with pKa 6.9 measured at 18 °C. Un-ionised H2S is much more toxic than the hydrosulfide ion, and it blocks cellular respiration in a similar way to cyanide. The toxic fraction therefore rises sharply as pH falls. Boyd's table gives un-ionised H2S as a percentage of total sulfide:

An oxidised layer at the sediment surface normally oxidises sulfide back to sulfate before it reaches the water. If that layer is lost, or sulfide arrives faster than it can be oxidised, a toxic residual concentration can build up in the water column.

What an ORP reading means chemically

Redox potential (Eh, also called ORP) is the tendency of a solution to gain or lose electrons, in volts or millivolts, relative to the standard hydrogen electrode at 0 V. Positive values indicate oxidising conditions and negative values reducing ones. Real measurements pair an inert platinum electrode with a reference electrode such as silver chloride or calomel, so the reading depends on the reference used. The ORP probe guide covers the hardware.

Boyd shows how little ORP responds to oxygen while any oxygen remains. Fresh water at pH 7 and 25 °C with 8 mg/L dissolved oxygen has a theoretical redox potential of 0.802 V against hydrogen. Against a calomel electrode (offset 0.242 V) that reads as 0.560 V, and at only 1 mg/L dissolved oxygen it still reads 0.547 V. To compare readings at pH 7, he adjusts by 0.0592 V per pH unit. Each 1 °C rise increases redox potential by about 0.0016 V. On this practical scale, where well-oxygenated water reads about 560 mV, denitrification begins near 300 mV. Nitrate is used up and iron and manganese oxides take over near 200 mV, and sulfate reduction and methane production start near 0 mV.

Natural water contains many redox couples that are not in equilibrium with each other, and electrodes respond slowly or become fouled, so measured potentials seldom match calculated values. Redox measurement is most useful for tracking change rather than reading an absolute value. A probe in the water column also cannot resolve millimetre-scale gradients in the substrate. For pond management, Boyd recommends measuring dissolved oxygen and watching sediment colour and odour instead.