Relevant To Their Interaction With Life Processes

Elements were formed in stars billions of years ago for the metals on Earth. Stars mainly burn hydrogen to helium, and later helium to carbon with oxygen and nitrogen as intermediates; higher elements are just "accidentally" formed. Thus, the probability for any element to be formed decreases hyperlogarithmically with its atomic number. The heavier any element is, the lower is the probability to find it in the Earth's crust, and the lower is the probability that evolution has used it. The three exceptions to this rule are [1]:

• The element synthesis in the stars stops at iron; thus, iron and its neighbors in the periodic system are present in high amounts on Earth despite their relatively high atomic masses.

• Li, Be, and B, which are not heavy metals, are overstepped during element synthesis; thus, they are present in low amounts on Earth, despite their relatively low atomic masses.

• Some heavy metals with extraordinarily stable nuclei were formed in higher amounts than any other heavy metals.

Metals with a density beyond 5 g/CC are heavy metals. Therefore, the term "heavy metals" is ascribed to the transition elements from V (but not Sc and Ti) to the half metal As; from Zr (but not Y) to Sb; and from La to Po, the lanthanides and the actinides. There are 21 nonmetals, 16 light metals, and 53 heavy metals (including As) constituting the total of 90 naturally occurring elements [10]. The transition elements with incompletely filled d-orbitals are heavy metals. The ability of the heavy metal cations to form complex compounds, which may or may not be redox active, is provided by the d-orbitals.

The origin of life is reduced to that of the biosphere, which from its inception was a complicated self-regulating system. A great variety of geochemical functions of living matter emerged at least as a result of the fact that any most primitive cell, being in aqueous and marine medium, had the closest possible contact with all the chemical elements of Mendeleev's table. Because life depends on water, heavy metals interesting for any living cell must form soluble ions. All the 53 heavy metals do not have a positive or adverse biological function because of the nonavailability of some heavy metals to the living cell in the usual ecosystem [11]. To enable the bioavailability, heavy metals must be present at least in nanomolar concentration in a given ecosystem because a concentration of 1 nM means that, in a cell suspension of 109/ml, each cell may receive about 600 ions. Metal ions generally present in lower concentration may be used by a microorganism for very specific purposes; however, the lower the mean concentration of the metal ion is in an ecosystem, the lower is the probability that any species carries around genes to use or detoxify this specific heavy metal [11].

The living conditions in the world ocean were most favorable, so it is possible that the Earth's hydrosphere was distinguished by the constancy of the biomass throughout the whole period of existence. Weast [10] has differentiated heavy metals into four classes on the basis of their concentration in seawater:

• Frequent possible trace elements with concentrations between 100 nM and 1 |M are Fe, Zn, and Mo.

• Possible trace elements with concentration between 10 and 100 nM are Ni, Cu, As, N, Mn, Sn, and U.

• Rare possible trace elements would be Co, Ce, Ag, and Sb.

• Cd, Cr, W, Ga, Zr, Th, Hg, and Pb are just below the 1-nM level.

Other elements, e.g., Au with 55.8 pM in seawater, are not likely to become trace elements.

The relative solubility of heavy metals under physiological conditions dictates the difference in biological importance and the toxicity of heavy metals vis-a-vis affinity to sulfur and other interaction with macrobioelements [11]. Because of the low solubility of the tri- or tetravalent cation [10], Sn, Ce, Ga, Zr, and Th have no biological influence. Of the remaining 17 heavy metals, Fe, Mo, and Mn are important trace elements with low toxicity; Zn, Ni, Cu, V, Co, W, and Cr are toxic elements with high to moderate importance as trace elements; and As, Ag, Sb, Cd, Hg, Pb, and U have no beneficial function, but must be considered by cells as toxins [11].

In addition to playing an important catalytic role by complexing with biochemicals, heavy metal cations in sophisticated biochemical reactions, such as nitrogen fixation; water cleavage during oxygenic photosynthesis; respiration with oxygen or nitrate; one-electron catalysis, rearrangement of C-C bonds; hydrogen assimilation; cleavage of urea; transcription of genes into mRNA, etc., can also form nonspecific complex compounds at higher concentrations in the cells that lead to the toxic effects. The heavy metal cations like Hg2+, Cd2+, and Ag+ are immensely toxic complex formers and are often perilous for any biological function. Even most important trace elements like Zn2+ or Ni2+ — and especially Cu2+ — are toxic at higher concentrations. Therefore, it was compelling to every life form to evolve some homoeostasis system for maintaining a tight control over intracellular concentration of heavy metal ions [11].


The divalent cations Mn2+, Fe2+, Co2+, Ni2+, Cu2+, and Zn2+ have ionic diameters in a range of 138 to 160 pm [10], a difference of 14%; all are, of course, double positively charged. Oxyanions like chromate, with four tetrahedrally arranged oxygen atoms, double negatively charged, differ mostly in the size of the central ion, so the structure of chromate resembles sulfate. The same is true for arsenate and phosphate. Thus, uptake systems for heavy metal ions must bind those ions tightly if they want to differentiate between a couple of similar ions. However, tight binding costs time and energy.


The uptake mechanisms that exist within a cell allow the entry of metal ions, including heavy metal ions that affect toxicity. There are two general uptake systems: the first one is driven by a

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