Lecture No. 15. Photosynthesis-Linked Periodic Variations; Trace Metal Toxicity.

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Reading Assignments:

1. Required reading: Fuller, Christopher C., and James A. Davis, 1989, Influence of coupling of sorption and photosynthetic processes on trace element cycles in natural waters, Nature 340 (6228), 52-54
2. Strongly recommended reading: Berner, Robert A., 1981, A new geochemical classification of sedimentary environments, Journal of Sedimentary Petrology 51 (2), 359-365

Announcements:

1. The term paper outlines were due Thursday, March 25. If you need more time, I will extend the deadline to Thursday, April 1. You are not expected to know everything you are going to put in your paper. However, by now you should have a reasonable idea of what the paper is about and how it is going to be organized. The outline should not be more than 1-2 pages.
2. The term papers will be due Thursday, April 29 rather than April 22.
3. In our next class meeting (Thursday, March 25) I will tie up a few loose ends. Any questions you may have about the midterm exam will be answered, and a few other things related to earlier lecture material will be given.
4. The field trip will be on Saturday, April 10. We will leave at 8:00 a.m. and return around 5:00 p.m. We will stop somewhere for lunch, probably in the Butte area.

Daily Cyclic Fluctuations in Water Chemistry

In their paper in Nature, Fuller and Davis discuss changes in stream geochemistry linked to the day-night cycle. The study site is a creek in a mining-impacted district in South Dakota. Whitewood Creek is a tributary of the Belle Fourche River on the northern fringe of the Black Hills. The groundwater passes through metal-contaminated floodplain sediments and enters the river in gaining reaches. This leads to the deposition of arsenic-enriched iron oxyhydroxides on the bottom of the stream. The pH of the creek water is between 8 and 9; it sometimes fluctuates by as much as 0.5 in any given 24 hour period.

Figure 1(a) from this paper shows the variations in pH over a period of about 60 hours during August 1987. The pH is higher when the sun is up and drops at sunset. The authors monitored stream pH about once an hour during this period and took hourly light readings with a radiometer during the daylight hours. They also periodically took water samples for analysis. During the period shown, streamwater pH was about 8.4 at noon and fell to values between 8.0 and 8.1 after dark. The pH variation was very nearly in phase with the light intensity readings. Variation of dissolved arsenate at two sites on the creek is shown in Fig. 1(b) and 1(c). The arsenic peaks and troughs lag the pH (and the sunlight cycle) by several hours.

The arsenic present in water and sediments was predominately arsenate — As (V). Arsenate has a strong affinity for iron oxyhydroxides, particularly amorphous or poorly crystallized ones like ferrihydrite. Arsenate becomes more ionized, and thus more negatively charged, at higher pH. Since ferrihydrite also is becoming more negatively charged with increased pH, the adsorption-desorption equilibrium is displaced to the left.

H₂AsO₄^– (aq) ó H₂AsO₄^– (adsorbed)

However, much of the arsenate is adsorbed to internal surfaces — surfaces of ferrihydrite that have become covered over by later-forming precipitates. It takes arsenate some time to work its way out into solution, since it has to diffuse along surfaces within aggregates of coagulated ferrihydrite particles. Thus the dissolved arsenic concentration is partially controlled by equilibrium and partly by kinetics.

The question arises: What is causing the pH fluctuations? The most obvious answer is the presence or absence of photosynthesis. Algae and other aquatic plants remove carbon dioxide and produce oxygen during daylight hours:

CO₂ + H₂O + hn ¨ CH₂O + O₂

All other things being equal (such as temperature, dissolved metal concentrations, nutrient levels, etc.), the rate of photosynthesis is directly proportional to the supply of sunlight. Carbon dioxide is removed by photosynthesis and is being replenished by diffusion of atmospheric carbon dioxide into the water. This produces a steady-state concentration of CO₂ somewhat lower than saturation. Because there is less carbonic acid dissolved in the water, the pH increases. At the same time, the production of oxygen by photosynthesis is opposed by diffusion of this oxygen out into the atmosphere and consumption of oxygen by aerobic organisms in the water and sediments. The net effect is a dissolved oxygen concentration higher than we would see without photosynthesis. The combined effect of all these processes is a higher level of dissolved oxygen and a higher pH than the water would have without photosynthesis.

When photosynthesis stops, a CO₂ sink and an O₂ source disappear, so the CO₂ concentration increases and the dissolved O₂ concentration decreases. The water becomes somewhat more acidic and somewhat less oxic. Because the authors are looking only at As (V) adsorption and desorption, they are concerned only with the pH changes and not the E_h changes.

The authors also performed a set of laboratory experiments using synthetic ferrihydrite (prepared from reagents) and artificially induced pH changes. These experiments largely confirmed what the field data showed: arsenic desorption took 6-8 hours to go about 50% to completion and still was increasing slowly after 24 hours. Of course, in a natural system in summer, the process reversed itself at sunrise, which is about 10-11 hours after sunset in August at the latitude of the study site.

The authors also noted that the arsenic showed less tendency to adsorb to the ferrihydrite at the end of the experiment. The ferrihydrite presumably was slowly converting to a more crystalline form (goethite?). The amplitude of the changes in dissolved arsenate concentration was larger in the lab than it was in the field. The authors suspect that there is some pH buffering in the bed sediments, so that porewater pH may be varying less than pH in the stream. They also suggest that the ferrihydrite in the creek sediments is more aged than the fresh ferrihydrite in the laboratory study and so has less sorption capacity and slower sorption-desorption kinetics.

Whitewood Creek is a fairly alkaline stream. The authors suggest that metal cycling in more acidic streams might be much larger than in an alkaline stream. They also suggest that iron oxyhydroxides might be cycled by redissolution and reprecipitation at a fairly high rate in some cases, which could explain the prevalence of poorly crystallized oxyhydroxides over more crystallized forms in many streams.

What constitutes a more acidic stream? How about the Clark Fork? It has a pH near 7. This is in the neighborhood of pH values that are right for desorption of many metals from sediments: zinc, copper, lead, etc.

Chris Brick sampled water in the Clark Fork near Deer Lodge over a period of about 48 hours as part of her dissertation research. Samples were taken and pH and dissolved oxygen were measured every couple of hours. She found that pH and dissolved oxygen followed a sinusoidal curve that was at a maximum around local noon and a minimum around midnight. (The maxima were offset somewhat from Mountain Daylight Time because true solar noon was about 1 ¸ hours later than noon MDT.) Dissolved oxygen and pH were in phase with each other. Dissolved zinc was about 180¡ out of phase; it hit a minimum at noon and a maximum around local midnight. This is what one would expect if zinc sorption and desorption were governed by pH.

Magenta = Zinc; Yellow = pH; Aqua = Dissolved Oxygen

Trace Metal Toxicity

Trace metals and metalloids are biologically important for several reasons.

1. They are required in some enzymes (biocatalysts) required for metabolism. E.g., selenium is required for glutathione peroxidase. Copper is necessary in enzymes that are used to make hemoglobin.
2. They are constituents of some vitamins. E.g., vitamin B-12 contains cobalt.
3. They are important in electron transfer. E.g., the Fe (II)–Fe (III) couple is what makes hemoglobin work. Other couples, such as Cu (I) – Cu (II) and Mo (V) – Mo (VI), are also used in essential processes.

Biologically there is a need for many trace elements. Oversupply and undersupply are both dangerous.

Thus there are two situations. Some elements are essential but become toxic at high levels. Other elements are not essential. They are tolerated at low levels and become toxic at high doses. Interestingly, the symptoms of the deficiency disease are often very similar to the symptoms of poisoning.

Here is the situation with essential elements:

Here is the situation with nonessential elements:

Toxicity is affected by certain characteristics: whether or not the elements are used by organisms, how available they are, and how abundant they are. The phase in which a metal is present is extremely important. Metals in the residual phase (e.g., silicates) are not very toxic, simply because they are unavailable. Metals in the exchangeable phase are much more available and therefore are more toxic.

Elements can be classified as

• Noncritical
• Toxic but insoluble or rare
• Very toxic and relatively accessible

See the handout. Note that many essential elements are listed as very toxic and relatively accessible. That is because they can be accessible at higher than optimal levels.

There are a number of factors that influence the toxicity of a heavy metal. These include

• The form of the metal in water — organic or inorganic, soluble or particulate, and the form the dissolved or particulate phase takes.
• The presence of other metals or other poisons. They may interact either positively (joint action, or more-than-additive effect) or negatively (antagonism, or less-than-additive effect), or they may have no interaction (strictly additive effects). For example, selenium is antagonistic toward arsenic, i.e., it lessens arsenic toxicity.
• Factors that influence the physiology of the affected organisms and/or the possible form of the element in water, such as temperature, pH, dissolved oxygen, light, or salinity.
• The condition of the organism: stage in life, age and size, sex, nutritional status, etc.
• Possible behavioral response of the organism.

The handout also lists tolerance levels for some metals and compounds in various species, including humans.

Toxicity Standards

I. Human Drinking Water Criteria

These are defined as concentrations above which water is considered unsafe to drink. These standards are relatively conservative. Some are based on epidemiological data. Here are some of the criteria for metals and metalloids. Those marked with an asterisk are under consideration for changes (usually reductions) in the US.

II. Aquatic Life Criteria

These criteria are of the form exp [a ln (hardness) + b], where hardness is in mg CaCO₃/L. Metal concentrations are expressed as total recoverable metals, mg/L. Metal toleration is usually better in hard water.

Here are some typical acute toxicity levels (mg/L) in the Clark Fork (150 mg/L hardness):

These are freshwater standards. There are also separate standards for seawater, most of which are more stringent.

To previous lecture: Lecture No. 14. Mining-Derived Metal Contamination (Continued)

To next lecture: Lecture No 16. Metal Toxicity II

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