Lecture No. 9. Classes of Geochemical Environments

Now to introduce a very useful system of classification for geochemical environments.

Berner’s Classification of Geochemical Environments

Berner starts by noting that there are two things we can measure easily down to about 1 m M: dissolved oxygen and hydrogen sulfide. These two species are very useful because:

• They are very strongly tied to redox reactions.
• Their concentrations have very major effects on organisms.
• They exert a major amount of control on the formation of authigenic minerals.
• They cannot coexist.
• Bacteria divide neatly into aerobes (killed by H₂S) and anaerobes (killed by oxygen). (Note: under this classification, facultative anaerobes are considered to be aerobes.)

Berner’s scheme:

• If dissolved oxygen is > 0.5% of saturation (~1 m M), the environment is oxic.
• If dissolved oxygen is < 0.5% of saturation, it is anoxic.

For an anoxic environment, if sulfide species are > 1 m M, it is sulfidic. If sulfide species are < 1 m M, it is postoxic or methanic.

For an oxic environment,

15 O₂ (aq) + 4 FeS₂ + 8 H₂O ¨ 2 Fe₂O₃ + 8 SO₄^2– + 16 H⁺

O₂ (aq) + 4 FeCO₃ ¨ 2 Fe₂O₃ + 4 CO₂

O₂ (aq) + MnS ¨ MnO₂ + S

In other words, oxygen + sulfides ¨ CO₂, SO₄^2–, and metal oxides.

For an anoxic-sulfidic environment,

3 H₂S (aq) + S + Fe₂O₃ ¨ 2 FeS₂ + 3 H₂O

(Note the complexity of this redox reaction! S starts out in 2 different redox states.)

H₂S (aq) + S + FeCO₃ ¨ FeS₂ + CO₂ + H₂O

2 H₂S (aq) + MnO₂ ¨ MnS + S + 2 H₂O

H₂S (aq) + MnO₂ + CO₂ ¨ MnCO₃ + S + H₂O

H₂S + oxides ¨ sulfides ± carbonates

Consider a water-sediment interface (see diagram on board), say at pH 7. If there is not much sulfur present, then Fe²⁺ predominates at low E_h (which is to say, a little way down into the sediments). If there are free Fe²⁺ and Mn²⁺ in the pore water, then the system is anoxic.

What drives their reactions?

"Organic matter" is approximately (CH₂O)₁₀₆(NH₃)₁₆(H₃PO₄)₁.

In oxic environments,

Organics + O₂ ¨ H₂O + CO₂ + HNO₃ + H₃PO₄.

(This is aerobic respiration as microbes do it.)

In sulfidic environments, we see sulfate as the oxidizing agent:

Organics + SO₄^2– ¨ CO₂ + NH₃ + H₃PO₄ + H₂S + H₂O

Typically,

Fe²⁺ + H₂S ¨ FeS + 2 H⁺,

followed by

FeS +S ¨ FeS₂

For iron, pyrite is the stable form. For other metals, one or another sulfide will form.

The iron came from the reduction of iron (III). We showed hematite reduction above, to simplify things. Actually, iron oxyhydroxides are more commonly the candidates for reduction.

2 FeOOH + H₂S ¨ S + 2 Fe²⁺ + 4 OH^–

For all this to happen, it is necessary to have

• enough S (usually present),
• enough bacteria (never a problem), and
• enough organic matter (food for the bacteria).

Anoxic Non-Sulfidic Environments

1. Low Organics (= Postoxic)

When oxygen is exhausted, microbes turn to other energy sources: first Mn (IV), then nitrate, then Fe (III).

Organics (low) + MnO₂ + H⁺ ¨ Mn²⁺ + CO₂ +H₃PO₄ + N₂ (or NH₃)

Organics (low) + NO₃^– ¨ CO₂ + H₃PO₄ + N₂ (or NH₃)

Organics (low) + Fe (III) ¨ CO₂ + H₃PO₄ + Fe²⁺ + N₂ (etc.)

Since organics are low, the oxidizing species may never be used up, and sulfate reduction does not get started.

2. Low Sulfur, High Organics (Methanic)

Sulfur was low to begin with or was exhausted, and there are plenty of organics left. This is typical of some freshwater systems such as marshes. Once oxygen, Mn (IV), nitrate, Fe (III), and sulfate are gone, fermentation and carbon dioxide reduction set in. Some species also reduce water!

Organics ¨ CO₂ + CH₄ + NH₃ + H₃PO₄ + H₂O + H₂

Carbohydrates + H₂O ¨ formic acid + H₂

4 H₂ + CO_2¨ CH₄ + 2 H₂O

Memorize the following table:

More on Particulate Matter

Metals and metalloids can be carried by sediment in several ways:

• as exchangeable metals,
• bound to carbonates,
• bound to Fe/Mn oxyhydroxides,
• bound to organic matter, or
• as part of crystalline minerals.

There are a number of extraction methods that are designed to separate out these different fractions. They are described in Elder’s article. Here is a brief synopsis of the extraction methods:

The various mechanisms by which metals bind to solids are sensitive to both pH and E_h. The pH dependence is due to several causes:

1. The mineral (or organic) surfaces carry a different electrical charge depending on the pH. The overhead shows the surface charges vs. pH for a number of mineral solids. The equivalence point is the pH at which the surface charge is zero. Note that FeOOH is negatively charged only over pH 7, but MnO₂ is negatively charged down to pH 4. Most metals are cations in solution; a cation will only be attracted to a negatively charged surface, and it will therefore only stick to the surface at pH values above the equivalence point for that solid.
2. At lower pH, there will be many more hydrogen ions in solution. Since they are also cations, H⁺ ions compete with other cations and may displace them from a negatively charged surface.
3. Some metals (e.g., Cr(VI), Mo(VI), W(VI), V(V)) exist as anions in solution. Arsenic and selenium also usually exist as anions. These species will be attracted to positively charged surfaces. However, this is complicated by the fact that low pH values tend to suppress their ionization. At high pH, on the other hand, there will be a higher concentration of hydroxide ions competing for adsorption sites. For some anions, calcite or g -Al₂O₃ might be the best adsorbent, since they carry positive charges at moderate pH.
4. Organic chelating agents generally work best for cations at higher pH, both because they are more soluble (acid precipitates humic acid) and because H⁺ competition is less.
5. Neutral compounds such as H₃AsO₃ are attracted to neutral or near-neutral substrates. However, the dipole moment or the polarizability of a molecule may be important. Arsenous acid is polar, even though it is uncharged at most pH values. It may therefore sorb to iron or manganese oxyhydroxides because a charged area on its molecule binds to an oppositely charged area on the substrate.
6. Specific interactions exist. Arsenic (V) tends to coprecipitate with FeOOH. Depending on the pH, it may form scorodite (FeAsO₄) or a sorbed arsenate salt. Arsenic (III) tends to become incorporated into authigenic pyrite. Barium readily forms carbonate or sulfate precipitates if the solution concentrations are high enough.

Another type of specific interaction is the greater or lesser affinity of a metal for organic matter. Here we are talking about humic acids (which are precipitated by acid), fulvic acids (which are not precipitated by acids, and which usually make up about 50% or the dissolved organics in surface waters) and other organics, such as the glycoproteins in bacterial slimes. Iron has a very high affinity for humic substances. On the other hand, manganese and cobalt do not have much affinity for humics. Copper (especially) and lead have high affinities for organics.

Organic matter tends to sorb strongly to clay minerals. Other ligands (say glutamic acid or picolinic acid) also sorb to mineral phases. With the lighter organics, you have a sort of crapshoot. If the molecule sorbs to the mineral phase in such a way that its chelating sites are hindered, it becomes less effective at capturing metals. If it sorbs in a position that leaves its chelating sites exposed, it will tend to take metals from solution to particulate phases.

Recall Horowitz’s assumption that metals are almost all carried on fine particles. Some metals avoid the fines! Examples are the alkaline and alkaline earth metals, such as Na, Sr, and Ba, which seem to have low affinities for clays, organics, and oxyhydroxides (which concentrate in the fines) and higher affinities for quartz and other (usually well-crystallized) minerals in the coarse fraction.

Elder gives another example of size-dependent metal sorption from the Saddle River in New Jersey. In the upstream region, zinc concentrated in the fines. Downstream the zinc distribution became bimodal, with a second peak for coarse particles. The coarse matter with an affinity for zinc was apparently floc from sewage plants.

Aquatic Biota

Both microorganisms and larger plants and animals have a number of effects on contaminant transport. Here are a few of those effects:

1. Microorganisms sorb metals to their outside surfaces and sometimes in the slime they secrete. (Plants and animals do too, but their surface areas are minuscule next to those of bacteria, microalgae, and fungi.)
2. Both plants and microorganisms take in metals and metalloids and sometimes alter them chemically. A common example is the methylation of arsenic by algae and cyanobacteria. Another is the methylation of tin by grasses of the genus Spartina. Sometimes the alteration is to an insoluble form that may be stored somewhere in the organism’s tissues or the bacterial cell. (This is a type of bioaccumulation.) This insoluble form may be re-released when the organism dies.
3. Biomagnification can lead to higher concentrations up the food chain. Mercury in fish muscle tissues is a famous example. Tuna is not high in mercury from industrial pollution; it is high because tuna are high on the food chain.
4. Some organisms secrete complexing agents (e.g., siderophores) that chelate metals.
5. There can be both reinforcement effects and antagonistic effects. An example of reinforcement given by Elder is the reinforcement of copper uptake by Mytilus edulis by high concentrations of lead, zinc, or silver. An example of antagonism is the reported effect of high zinc in retarding lead uptake by humans. Another example is the effect of high potassium in suppressing cesium uptake by plants. A third example is the retarding effect calcium has on heavy-metal uptake by fish.
6. Photosynthetic plants (including microalgae) modify the pH and E_h in water. Their uptake of CO₂ raises pH by removing an acid, while the oxygen they release makes the water more oxic. It has been reported that river pH can drop by several tenths or a unit after sunset, while oxygen levels may drop enough to allow reduction of MnO₂ coatings and release of adsorbed metals.
7. Benthic invertebrates can kick up a lot of sediment during the times when they are active. If the bugs come out from under rocks after sunset, then the suspended load may increase at the same time the pH drops. Spawning fish can create a lot of bioturbation seasonally.

To previous lecture: Lecture No 8. Eh-pH Diagrams II

To next lecture: Lecture No 10. Some Handy Computational Tools

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