The Sulfur Cycle

Steven Dutch, Natural and Applied Sciences, University of Wisconsin - Green Bay
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Global Sulfur pools

Almost all sulfur is stored in the crust, with minor amounts in sea water, atmosphere, biosphere

Global Sulfur emissions to land:
(1 Teragram=1012 grams, or approximately a million tons)

Source  Size (in Tg/year)

Volcanoes  12
Sea Spray  300
Marsh and Marine Decomposition  5
Di-Methyl Sulfide (DMS) produiction  39
Volitization from land 64
Wind erosion 20
Total natural emssions 440

Human fossil fuel combustion in the Northern Hemisphere  120

Thus, approximately 20% of the total sulfur emssions on the planet are due to the people of North America and Eurasia!

Biological uses of Sulfur


Sulfur is not one of the major nutrients required for life.  However, it is necessary is small amounts.

Sulfur is used to to create di-sulfide bonds in proteins and enzymes.  These bonds help hold these molecules in the shapes which their functions require.  Without Sulfur, these molecules would not function.

Effect of excess sulfur emissions on the biosphere


In an anerobic environment, Sulfur will combine with Hydrogen to form Hydrogen Sulfide (H2S), which smells like rotten eggs.  This creates the bad smell characteristic of wetland areas.

When Sulfur is burned in an aerobic environment, it creates a number of different sulfur oxides (SO, SO2).  These compounds will eventually be converted into SO3 by sunlight.  SO3 can combine with water to form Sulfuric Acid:

SO3 + H2O = H2SO4

This strong acid will dissolve in cloud and rain droplets, and then be carried to the ground.


Beacuse of Carbonic Acid, rainwater has always been slightly acidic (pH=5.6).  But, this is a weak acid.  Adding sulfuric acid into rain causes the pH to fall greatly (around 4.6 in E. Wisconsin to 4.2 in the northeast).

As the pH scale is logarithmic, this means the rainwater of eastern Wisconsin is now 10 times more acidic than normal, with areas in the northeast being over 50 times more acidic.

The increased acidity of rainwater will often lead to increased acidity of surface and groundwater.  This is not as bad of a problem in areas underlain by Calcium-rich rocks, as the calcium will interact with carbonic acid to form Calcium Bicarbonate, which is the antacid found in Tums:

Ca+2 + 2H2CO3- = Ca(HCO3)2 + 2H+

In areas underlain by granite, the acid will not be counteracted because granite is poor in elements like calcium and magnesium that can neutralize acids, and the lake acidity will rise to the level of the rainwater.


This is important, as at high water acidity, many fish will not reproduce and soil leaching will be greater, washing metal ions (such as Aluminum) into surface and ground water.  This may also impair the ability of aquatic species to live.

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Created 2 September 2011, Last Update 02 September 2011

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