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The Anatomy Of The Respiratory System Biology Essay

There is a need to distinguish between breathing and respiration because these two terms are confused greatly. Breathing and respiration are NOT the same, but because of their interrelationship, people often refer to both terms as one. However, breathing is a simple mechanical process where air enters and exits the lungs and is due to volume and pressure changes. Respiration on the other hand, is a more complex process, and it involves generally a twofold process: (i) delivering oxygen to the tissues for cell metabolism and (ii) removing carbon dioxide. This assignment deals with respiration as it relates to Henry’s Law. But before the relation is explained, there is a need to briefly explain respiration (and by extension gas exchange) and how it relates to the body, then it will be easier to see the relationship between the law.

The respiratory system may be separated into two different zones.  These zones are called the conductive zone and the respiratory zone. The conductive zone deals with pulmonary ventilation, or the movement of air in and out of the lungs. The respiratory zone deals with external respiration, or the exchange/diffusion of O2 and CO2 between the alveoli and blood. Upon completion of external respiration, O2 is transported via the cardiovascular system to the tissue where internal respiration removes O2 from the hemoglobin and transports the O2 to the cell for cellular respiration. 

Figure 1 – the diagram shows simply how O2 and CO2 is exchanged. In the lungs oxygen and carbon dioxide (a waste product of body processes) are exchanged in the tiny air sacs (alveoli) at the end of the bronchial tubes. The alveoli are surrounded by capillaries. When a person inhales, oxygen moves from the alveoli to the surrounding capillaries and into the bloodstream. At the same time, carbon dioxide moves from the bloodstream to the capillaries and into the alveoli. The carbon dioxide is removed from the lungs when a person exhales. (Rhoads, 2008)

Figure 2 - The image at shows blood being oxygenated at the lungs (red) and traveling through the heart to the systemic circuit.

The gas exchange at the alveoli/blood interface is called external respiration. Gas exchange between blood and tissue cells is termed internal respiration. Deoxygenated blood (blue) flows from the tissues in the systemic circuit back to the heart, the lungs, and CO2 is exhaled. As seen in the illustration, the cardiovascular system (the heart and blood vessels) plays a mechanical role in transporting gases, but partial pressures are also involved.

VAPOUR PRESSURE

When a liquid is exposed to open space above its surface, molecules of dissolved gas will escape from the liquid surface into the space above. Equilibrium is reached when the rate of escape of molecules from the liquid phase equals the rate at which gas molecules reenter the liquid phase. The vapour pressure of a gas over a liquid is the partial pressure exerted by gas molecules when there is an equilibrium between the liquid and gas phases for these molecules. In repiratory physiology, the most common vapour pressure measured is that for water. At equilibrium, the partial pressure of water in the gas phase is equal to its vapour pressure. The vapour pressure of water is solely dependent on temperature. (Hlastala, 2001)

The solubility of gases in liquids is important to our understanding of gas transport mechanisms. The amount of gas that can dissolve in a liquid is directly proportional to the partial pressure of that gas above the liquid. This is a statement of an important physical law termed Henry’s law. It is important to distinguish the amount of gas dissolved in the liquid from that which may chemically combine with the liquid. Henry’s law refers only to the quantity of gas that is dissolved at equilibrium, when the partial pressures of a given species in the gas and liquid phase are equal by the following relationship: (Hlastala, 2001)

Where:

denotes a single gas species.

represents the concentration of dissolved species now in the liquid phase.

Typical units for this quantity (of gases dissolved in blood) are milliliters (mL) of dissolved gas at STPD per deciliter of blood (also termed volume percent). The proportional factor, (βx, is the solubility coefficient. Its units are milliliters of dissolved gas at STPD per deciliter of blood per mmHg. The solubility coefficient depends on temperature; the higher the temperature, the smaller its value. Thus, at a higher temperature, less gas will dissolve in a given volume of liquid. (Hlastala, 2001)

The following example illustrates the use of Henry’s law to calculate the concentration of dissolved O2 in plasma when the surface of the plasma is exposed to fully saturated air. At 20 oC, the O2 solubility coefficient is 0.00364 mL O2 dl-1 mmHg-1 and take PO2 as 150 mmHg; therefore,

DALTON’S LAW OF PARTIAL PRESSURES

The total pressure exerted by a mixture of gases is the sum of the pressures exerted independently by each gas in the mixture. The pressure exerted by each gas, its partial pressure, is directly proportional to its percentage in the total gas mixture.

The air you breathe is a mixture of gases, composed primarily of nitrogen (N2), oxygen (O2), carbon dioxide (CO2), and water vapor (H2O). Each gas exerts a pressure as if the other gases were not present, and these partial pressures (written Pgas) can be summed to determine the total pressure (760 mm Hg at sea level).

As the amounts of gases in a mixture vary, their partial pressures will vary, as they are ratios of the total pressure. Thus, if you add oxygen to a mixture, its partial pressure will increase (given a constant volume). If you were to increase or decrease the pressure on this system, the partial pressures of the gases would increase or decrease, respectively, as well.

GAS

ATMOSPHERE (at sea level)

ALVEOLI (at sea level)

PARTIAL PRESSURE

(mm Hg)

PERCENTAGE

PARTIAL PRESSURE

(mm Hg)

PERCENTAGE

N2

596

78.4

569

74.9

O2

160

21.05

104

13.7

CO2

0.3

0.04

40

5.2

H2O

3.7

0.49

47

6.2

TOTALS

760 mm Hg

100.0%

760 mm Hg

100.0%

Gases flow from areas of higher pressure to those of lower pressures by diffusion. This is how smells migrate across a room. Looking at the table, it can be seen that nitrogen and oxygen gases will rush into the lungs and carbon dioxide and water vapor will rush out. There exist steep partial pressure gradients between the atmosphere and the lung alveoli which insure this. Because blood is a liquid and air is composed of gases, Henry’s law play an important role.

HENRY’S LAW

Henry’s law is written as the following expression:

Where:

is the mole fraction of the solute and

is an empirical constant (with dimensions of pressure).

is the partial pressure above the solute.

The amount of gas that will dissolve in a liquid varies directly with the pressure above that liquid (Hein, 1997). Or, when a mixture of gases is in contact with a liquid, each gas will dissolve in the liquid in proportion to its partial pressure (Marieb, 1992).

High pressures force gas into solution/dissolution. However, solubilities and temperatures also come into play when considering Henry's law. Even though a huge PN2 gradient may exist between the air and plasma, the fact is, nitrogen is barely soluble at all. (Caroline, 2000)

Consider Figure 3. Find the sad face. Here the blood is CO2 rich (45 mm Hg) and low in oxygen (40 mm Hg). This sad blood is being delivered to the lungs, where the alveolar PO2 is 104 mm Hg and the PCO2 is 40 mm Hg. See? You probably know right away that fresh oxygen will rush across the respiratory membrane into the blood and refresh its supply to 104 mm Hg. But the gradient of PCO2 is only 5 mm Hg. It would seem then that carbon dioxide would have a lesser tendency to evacuate the blood, but it turns out that carbon dioxide is about 20 times more soluble in blood and alveolar fluid than oxygen (Marieb, 1992). Even though its gradient is smaller, its solubility encourages its movement through the membrane into the alveoli. (Caroline, 2000)

The fresh, red blood is pumped by the heart to tissues throughout the body. Traveling down the systemic circuit, from aorta to arteries to arterioles to capillaries, oxygen is "used" and carbon dioxide is "created." Internal respiration at the tissues creates a new equilibrium of partial pressures of oxygen and carbon dioxide, so that again the venous (return) blood is in the same sad condition we started in. (Caroline, 2000)

In summary, arterial blood (red) has a higher PO2 and a lower PCO2 than venous blood, and this system is perpetuated at both ends by external and internal respiration. (Caroline, 2000)

Figure 3 - The image at shows blood being oxygenated at the lungs (red) and traveling through the heart to the systemic circuit. Also shows the partial pressures. (Caroline, 2000)

PRACTICAL CONSEQUENCES

We would reason that a deep diver is subjected to increased "atmospheric" water pressure, a great force pressing in on his chest. Remember, water is heavier than air and exerts much more pressure? To give an idea, atmospheric pressure increases by 760 mm Hg every 30 feet below sea level in sea water. At 2 feet, the atmospheric pressure will have increased 25 mm Hg to 785 mm Hg. It is generally accepted that below 2 feet, a diver breathing through a tube (a snorkel) will have extreme difficulty overcoming the weight of the water above him. Providing the air tube does not collapse, it provides atmospheric air at a lower pressure than the pressure pressing in on the thorax. At resting (where an equilibrium should occur), the lungs would have to expand to great proportions to equalize outside pressure. The little guy would try to take a breath and the great pressure would prevent his lungs from expanding to this size. If he was lucky at all, he might experience "rapture of the deep" for a few moments, but more likely he would suffocate sooner. (Caroline, 2000)

Figure 4 - The image shows that for every 33 feet you descend under sea water, atmospheric pressure increases by 1 atmosphere (760 mm Hg).

The diagram in Figure 4 shows that for every 33 feet you descend under sea water, atmospheric pressure increases by 1 atmosphere (760 mm Hg). This is a serious gradient, and detrimental to life. Finally, after enough people popped underwater, someone invented scuba (self-contained underwater breathing apparatus). Compressed air in a tank equalizes with the water pressure so the lungs can expand against the crushing force of deep water. Scuba allows people to explore depths for longer periods of time, but it also presents serious potential problems, namely decompression sickness (also called Caisson's disease and "the bends") and air embolism (Marieb, 1997; Caroline, 2000).

THE BENDS

Imagine you are scuba diving. The tank on your back is small but weighs a ton, probably because enough gas to fill a phone booth is compressed inside it, at about 3000 psi. As you descend, the pressure and amount of gas entering your lungs is harmlessly increased. Keeping the law of partial pressures in mind, you know that the partial pressures of all atmospheric gases entering your lungs are increased across the board. Henry's law explains why more gas will be forced into solution in the body tissues. Nitrogen levels are of concern. Nitrogen dwindles in lipid-rich tissues like the central nervous system (brain/spinal cord), bone marrow, fat, and dissoved in blood, and at prolonged concentrated levels causes "rapture of the deep," a sort of drunken stupor. (Caroline, 2000)

You've been under 100 feet for an hour already, and have just realized that you've been staring at the same coral for several minutes. You are suffering nitrogen narcosis. Panicking, you begin a hasty ascent. You've forgotten to go slow, to gradually decompress your tank. The partial pressure of nitrogen drops swiftly, and a sharp gradient forms between nitrogen in the bloodstream (lowered) and nitrogen dissolved in fatty tissue (high). Nitrogen "boils" out of the tissues, creating emboli (bubbles) in bones, joints, and muscle. The pain seems to shadow the itching, coughing, skin rash, and seizures. At the surface, you scream, you bob in the water for a few moments, and quickly you drop into shock. (Caroline, 2000)

One big rule of scuba diving is to "keep breathing." As chemistry whizzes, we know why: Boyle's law. At one depth, the pressure in the lungs will be appropriate for the ambient pressure. If one holds the breath and ascends, the pressure in the lungs will still be high, but the ambient pressure will drop. The lungs will expand, and alveoli may pop. If an alveoli ruptures in such a way that an open bridge is formed between outside air and the bloodstream, when a breath is taken at surface, air emboli bubble into circulation. Seizures, localized motor and sensory deficits, heart attack, and unconsciousness are likely results (Marieb, 1997; Caroline, 2000). It follows naturally that many people drown when unconscious in water. (Caroline, 2000)

POLLUTION

Pollution is the introduction of contaminants into a natural environment that causes instability, disorder, harm or discomfort to the ecosystem i.e. physical systems or living organisms

The major forms of pollution are listed below along with the particular pollutants relevant to each of them:

Air pollution, the release of chemicals and particulates into the atmosphere. Common gaseous air pollutants include carbon monoxide, sulfur dioxide, chlorofluorocarbons (CFCs) and nitrogen oxides produced by industry and motor vehicles. Photochemical ozone and smog are created as nitrogen oxides and hydrocarbons react to sunlight. Particulate matter, or fine dust is characterized by their micrometre size PM10 to PM2.5.

Water pollution, by the release of waste products and contaminants into surface runoff into river drainage systems, leaching into groundwater, liquid spills, wastewater discharges, eutrophication and littering.

CHEMODYNAMICS

Environmental chemodynamics is concerned with how chemicals move and change in the environment. There is a need to consider three specific partitioning relationships that control the ‘leaving’ and ‘gaining’ of pollutants among compartments, especially air, particles, surfaces and organic tissues. (Vallero, 2008). These concepts may be applied to estimating and modeling where a contaminant will go after it is released. These relationships are sorption, solubility, volatilization and organic carbon-water partitioning, which are respectively expressed by coefficients of sorption (distribution coefficient, KD, or solid-water partition coefficient, Kp), dissolution or solubility coefficients, air-water partitioning (and the Henry’s Law, KH, constant), and organic-water (KOC) partitioning. (Vallero, 2008)

In chemodynamics, the environment is subdivided into finite components. Thermodynamically, the mass of the contaminant entering and the mass leaving a control volume must be balanced by what remains within the control volume (ala the conversation laws). Likewise, within that control volume, each compartment may be a gainer or loser of the contaminant mass, but the overall mass must blanace. The generally inclusive term for these compartmental changes is known as fugacity or the ‘feeling potential’ of a substance. It is the propensity of a chemical to escape from one type of environmental compartment to another. Combining the relationships between and among all of the partitioning terms is one means of modeling chemical transport in the environment. This is accomplished by using thermodynamic principles and hence, fugacity is a thermodynamic term. (Vallero 2008; Harrison 2001)

The simplest chemodynamic approach addresses each compartment where a contaminant is found in discrete phases of air, water, soil, sediment and biota as shown in Figure 5.

Figure 5 - Relationship between airwater partitioning and octanolwater partitioning and affinity of classes of contaminants for certain environmental compartments. (Vallero, 2008)

However, a complicating factor in environmental chemodynamics is that even within a single compartment, a contaminant may exist in various phases (e.g. dissolved in water and sorbed to a particle in the solid phase). The physical interactions of the contaminant at the interface between each compartment are a determining factor in the fate of the pollutant. Within a compartment, a contaminant may remain unchanged (at least during the designated study period), or it may move physically, or it may be transformed chemically into another substance. Actually, in many cases all three mechanisms will take place. A mass fraction will remain unmoved and unchanged. Another fraction remains unchanged but is transported to a different compartment. Another fraction becomes chemically transformed with all remaining products staying in the compartment where they were generated. And, a fraction of the original contaminant is transformed and then moved to another compartment. So, upon release from a source, the contaminant moves as a result of thermodynamics. (Vallero 2008; Harrison 2001)

Fugacity requires that at least two phases must be in contact with the contaminant. For example, the Kow value is an indication of a compounds likelihood to exist in the organic versus aqueous phase. This means that if a substance is dissolved in water and the water comes into contact with another substance, e.g. octanol, the substance will have a tendency to move from the water to the octanol. Its octanolwater partitioning coefficient reflects just how much of the substance will move until the aqueous and organic solvents (phases) will reach equilibrium. So, for example, in a spill of equal amounts of the polychlorinated biphenyl (PCB), decachlorobiphenyl (logK of 8.23), and the pesticide chlordane (logK of 2.78), the PCB has much greater affinity for the organic phases than does the chlordane (more than five orders of magnitude). This does not mean that a great amount of either of the compounds is likely to stay in the water column, since they are both hydrophobic, but it does mean that they will vary in the time and mass of each contaminant moving between phases. The rate (kinetics) is different, so the time it takes for the PCB and chlordane to reach equilibrium will be different. This can be visualized by plotting the concentration of each compound with time (Figure 6). (Vallero 2008; Harrison 2001)

Figure 6 - Relative concentrations of a PCB and chlordane in octanol with time (Vallero, 2008)

When the concentrations plateau, the compounds are at equilibrium with their phase. When phases contact one another, a contaminant will escape from one to another until the contaminant reaches equilibrium among the phases that are in contact with one another. Kinetics takes place until equilibrium is achieved. (Vallero 2008; Harrison 2001)

PARTITIONING TO THE GAS PHASE: VOLATILIZATION

In its simplest connotation, volatilization is a function of the concentration of a contaminant in solution and the contaminants partial pressure. Henrys law states that the concentration of a dissolved gas is directly pro-portional to the partial pressure of that gas above the solution:

Where:

KH is the Henrys law constant; p is the partial pressure of the gas;

pa is the partial pressure of the gas;

[C] is the molar concentration of the gas

or,

Where:

CW is the concentration of gas in water.

So, for any chemical contaminant we can establish proportionality between the solubility and vapor pressure. Henrys law is an expression of this proportionality between the concentration of a dissolved contaminant and its partial pressure in the headspace (including the open atmosphere) at equilibrium. A dimensionless version of the partitioning is similar to that of sorption, except that instead of the partitioning between solid and water phases, it is between the air and water phases (K ): (Vallero 2008; Harrison 2001)

Where:

CA is the concentration of gas A in the air.

The relationship between the air/water partition coefficient and Henrys law constant for a substance is:

Where:

R is the gas constant (8.21 * 10-2 L at mol-1 K-1)

T is the temperature (K)

Henrys law relationships work well for most environmental conditions. It represents a limiting factor for systems where a substances partial pressure is approaching zero. At very high partial pressures (e.g. 30Pa) or at very high contaminant concentrations (e.g. 1000ppm), Henrys law assumptions cannot be met. Such vapor pressures and concentrations are seldom seen in ambient environmental situations, but may be seen in industrial and other source situations. Thus, in modeling and estimating the tendency for a substances release in vapor form, Henrys law is a good metric and is often used in compartmental transport models to indicate the fugacity from the water to the atmosphere. (Vallero 2008; Harrison 2001)

HENRY’S LAW EXAMPLE

At 25C, the log Henrys Law constant (logKH ) for 1,2-dimethylbenzene (C8H10) is 0.71 L atm mol-1 and the log octanol-water coefficient (log Kow) is 3.12. The logKH for the pesticide parathion (C10H14NO5PS) is -3.42 L atm mol-1, but its log Kow is 3.81. Explain how these substances can have similar values for octanol-water partitioning yet so different Henrys law constants. What principle physicochemical properties account for much of this difference? (Vallero 2008; Harrison 2001)

Both dimethylbenzene and parathion have an affinity for the organicphase compared to the aqueous phase. Since Henrys law constants are a function of both vapor pressure (Po) and water solubility, and both compounds have similar octanol-water coefficients, the difference in the Henrys law characteristics must be mainly attributable to the compounds respective water solubilities, their vapor pressures, or both. (Vallero 2008; Harrison 2001)

Parathion is considered semivolatile because its vapor pressure at 20C is only 1.3 * 10-3 Parathion’s solubility in water is 12.4 mg/L at 25oC. (Vallero 2008; Harrison 2001)

1,2-Dimethylbenzene is also known as ortho-xylene (o-xylene). The xylenes are simply benzenes with more two methyl groups. The xylenes have very high vapor pressures of 4.5 * 102 kPa, and water solubilities of about 200mgL at 25oC. (Vallero 2008; Harrison 2001)

Thus, since both the solubilities are relatively low, it appears that the difference in vapor pressures is responsible for the large difference in the Henry law constants, i.e. the much larger tendency of the xylene to leave the water and enter the atmosphere. Some of this tendency may result from the higher molecular weight of the parathion, but is also attributable to the additional functional groups on the parathion benzene than the two methyl groups on the xylene (Figure 7). (Vallero 2008; Harrison 2001)

Figure 7 - Molecular structure of the pesticide parathion and the solvents ortho-xylene, toluene, and benzene (Vallero, 2008)

Another way to look at the chemical structures is to see them as the result of adding increasing complex functional groups, i.e., moving from the unsubstituted benzene to the single methylated benzene (toluene) to o-xylene to parathion. The substitutions result in respective progressively decreasing vapor pressures:

Benzene’s Po at 20oC = 12.7 kPa

Toluene’s Po at 20 oC = 3.7 kPa

o-Xylenes Po at 20oC = 0.9 kPa

Parathion’s Po at 20oC = 1.3 * 10-3 kPa

The effect of these functional group additions on vapor pressure is even more obvious when seen graphically (Figure 8).

Figure 8 - Effect of functional group substitutions on vapor pressure of four organic aromatic compounds (Vallero, 2008)

It is important to keep in mind that Henrys law constants are highly dependent on temperature, since both vapor pressure and solubility are also temperature dependent. So, when using published KH, one must compare them isothermically. Also, when combining different partitioning coefficients in a model or study, it is important to either use only values derived at the same temperature (e.g. sorption, solubility, and volatilization all at 20 oC), or to adjust them accordingly. A general adjustment is an increase of a factor of 2 in KH for each 8oC temperature increase. (Vallero 2008; Harrison 2001)

Also, any sorbed or otherwise bound fraction of the contaminant will not exert a partial pressure, so this fraction should not be included in calculations of partitioning from water to air. For example, it is important to differentiate between the mass of the contaminant in solution (available for the KAW calculation) and that in the suspended solids (unavailable for KAW calculation). This is crucial for many hydrophobic organic contaminants, where they are most likely not to be dissolved in the water column (except as co-solutes), with the largest mass fraction in the water column being sorbed to particles. (Vallero 2008; Harrison 2001)

The relationship between KH and Kow is also important. It is often used to estimate the environmental persistence, as reflected in the chemical half-life ( ) of a contaminant. However, many other variables determine the actual persistence of a compound after its release. Note in the table, for example, that benzene and chloroform have nearly identical values of KH and Kow, yet benzene is far less persistent in the environment. (Vallero 2008; Harrison 2001)

However, relative affinity for a substance to reside in air and water can be used to estimate the potential for the substance to partition not only between water and air, but more generally between the atmosphere and biosphere, especially when considering the long-range transport of contaminants (e.g. across continents and oceans). Such long range transport estimates make use of both atmospheric T½ and KH. The relationship between octanol-water and air-water coefficients can also be an important part of predicting a contaminants transport. For example, Figure 5 provides some general classifications according to various substances Kow and KH relationships. In general, chemicals in the upper left hand group have a great affinity for the atmosphere, so unless there are contravening factors, this is where to look for them. Conversely, substances with relatively low KH and Kow values are less likely to be transported long distance in the air. (Vallero 2008; Harrison 2001)

Partitioning to Organic Tissue: Relatively hydrophobic substances frequently have a strong affinity for fatty tissues (i.e. those containing high Kow compounds). Therefore, such contaminants can be sequestered and can accumulate in organisms. In other words, certain chemicals are very bioavailable to organisms that may readily take them up from the other compartments. Bioavailability is an expression of the fraction of the total mass of a compound present in a compartment that has the potential of being absorbed by the organism. Bioaccumulation is the process of uptake into an organism from the abiotic compartments. Bioconcentration is the concentration of the pollutant within an organism above levels found in the compartment in which the organism lives. So, for a fish to bioaccumulate DDT, the levels found in the total fish or in certain organs (e.g. the liver) will be elevated above the levels measured in the ambient environment. In fact, DDT is known to bioconcentrate many orders of magnitude in fish. A surface water DDT concentration of 100ppt in water has been associated with 10ppm in certain fish species (a concentration of 10000 times!). Thus the straightforward equation for the bioconcentration factor (BCF) is the quotient of the concentration of the contaminant in the organism and the concentration of the contaminant in the host compartment. So, for a fish living in water, the BCF is: (Vallero 2008; Harrison 2001)

The BCF is applied to an individual organism that represents a genus or some other taxonomical group. However, considering the whole food chain and trophic transfer processes, in which a compound builds up as a result of predator/prey relationships, the term biomagnification is used. Some compounds that may not appreciably bioconcentrate within lower trophic state organisms may still become highly concentrated. For example, even though plankton have a small BCF (e.g. 10), if subsequently higher-order organisms sequester the contaminant at a higher rate, by the time top predators (e.g. alligators, sharks, panthers, and humans) may suffer from the continuum of biomagnification, with levels many orders of magnitude higher than what is found in the abiotic compartments. (Vallero 2008; Harrison 2001)

For a substance to bioaccumulate, bioconcentrate, and biomagnify, it must be at least somewhat persistent. If an organisms metabolic and detoxification processes are able to degrade the compound readily, it will not be present (at least in high concentrations) in the organisms tissues. However, if organisms endogenous processes degrade a compound into a chemical species that is itself persistent, the metabolite or degradation product will bioaccumulate, and may bioconcentrate, and biomagnify. Finally, cleansing or depuration will occur if the organism that has accumulated a contaminant enters an abiotic environment that no longer contains the contaminant. However, some tissues have such strong affinities for certain contaminants that the persistence within the organism will remain long after the source of the contaminant is removed. For example, the piscivorous birds, such as the Common Loon (Gavia immer), decrease the concentrations of the metal mercury (Hg) in their bodies by translocating the metal to feathers and eggs. So, every time the birds molt or lay eggs they undergo Hg depuration. Unfortunately, when the birds continue to ingest mercury that has bioaccumulated in their prey (fish), they often have a net increase in tissue Hg concentrations because the bioaccumulation rate exceeds the depuration rate. (Vallero 2008; Harrison 2001)

Bioconcentration can vary considerably in the environment. The degree to which a contaminant builds up in an ecosystem, especially in biota and sediments, is related to the compounds persistence. For example, a highly persistent compound often possesses chemical structures that are also conducive to sequestration by fauna. Such compounds are generally quite often lipophilic, have high Kvalues, and usually low vapor pressures. This means that they may bind to the organic molecules in living tissues and may resist elimination and metabolic process, so that they build up over time. However, the bioaccumulation and bioconcentration can vary considerably, both among biota and within the same species of biota. For example, the pesticide mirex has been shown to exhibit BCFs of 2600 and 51400 in pink shrimp and fathead minnows, respectively. The pesticide endrin has shown an even larger interspecies variability in BCF values, with factors ranging from 14 to 18000 recorded in fish after continuous exposure. Intraspecies BCF ranges may also be high, e.g., oysters exposed to very low concentrations of the organometallic compound, tributyl tin, exhibit BCF values ranging from 1000 to 6000. (Vallero 2008; Harrison 2001)

Even the same compound in a single medium, e.g. a lakes water column or sediment, will show large BCF variability among species of fauna in that compartment. An example is the so-called dirty dozen compounds. This is a group of persistent organic pollutants (POPs) that have been largely banned, some for decades, but that are still found in environmental samples throughout the world. As might be expected from their partitioning coefficients, they have concentrated in sediment and biota. The worst combination of factors is when a compound is persistent in the environment, builds up in organic tissues, and is toxic. Such compounds are referred to as persistent bioaccumulating toxic substances (PBTs). (Vallero 2008; Harrison 2001)

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