Compounds Of Metallic Elements

Metallic elements and metalloids form compounds in various oxidation states, yielding inorganic compounds such as salts and saltlike products, metal complexes (coordination compounds), and organometallic compounds. In metallic compounds, the atoms bind either in ionic or covalent bonds. Intermediate forms are also seen. Dissolved in water the metallic compounds dissociate into metal ions, mostly as cations. In some cases such as the permanganate ion (MnO4-), an oxoanion is formed. Metallic ions can form compounds with other metallic ions, forming alloys with two or more metals in varying proportions. Binary and multicomponent systems also exist in the crystalline phase (FeS is an example).

3.1 Covalent and Ionic Bonds

Two major types of chemical bonds exist (i.e., covalent and ionic). The covalent bond is defined as a region of relatively high electron density between nuclei that arises, at least partly, from the sharing of electrons and produces an attractive force and characteristic internuclear distance (McNaught and Wilkinson, 1997). Covalent bonds exist as homopolar and heterocova-lent. When the two atoms of the diatomic molecule are the same (e.g., two hydrogen atoms), the electron density is distributed symmetrically between the two nuclei, and the covalent bond is homopolar. If the two atoms are not the same, the electron distribution will be asymmetrical, and the electron density will be displaced toward the atomic nucleus that is more electronegative (i.e., which has a higher capacity to attract electrons). This is a heteropolar covalent bond. The greater the difference in electronegativity of two atoms forming a bond, the more uneven the distribution of the electrons will be. In an extreme case, a complete transfer of electrons from one atom to another occurs, thus forming an ionic bond. Metallic elements have low electronegativity. Chemical bonds are rarely entirely covalent or entirely ionic. Ionic bonds are predominantly formed in metal salts like chlorides (NaCl) or nitrates (Ca(NO3)2). Covalent bonds are predominantly, but not exclusively, formed between metals and carbon atoms as in organometallic compounds such as dimethylmercury (CH3-Hg-CH3).

3.2 Oxidation Number

Oxidation can be defined according to the following three criteria (McNaught and Wilkinson, 1997).

1. Oxidation is the complete removal of one or more electrons from a molecular entity (also called "de-electronation"), for example, the Zn2+ ion derives from the atom Zn of which 2 electrons have been removed.

2. This definition can be extended to chemical reactions in which a complete electron transfer does not occur and which, by custom and in current usage, are called oxidations. In this application, oxidation numbers are considered. Oxidation now is an increase in the oxidation number of any atom within a substrate, for example,

Fe2+ - e- o Fe3+. The oxidation number in an ion or a molecule is the charge the atom would have if the polyatomic ion or molecule was composed entirely of ions. Thus, for example, in MnO4-, manganese is considered to be in the oxidation state +7 (MnVI1), and oxygen is assumed to exist as O2- ion.

3. Oxidation is also the gain of oxygen and/or loss of hydrogen of an organic substrate. All oxidations meet criteria 1 and 2, and many meet criterion 3, but this is not always easy to demonstrate.

Alternately, an oxidation can be described as a transformation of an organic substrate that can be rationally dissected into steps or primitive changes. The latter consist in removal of one or several electrons from the substrate followed or preceded by gain or loss of water and/or hydrons or hydroxide ions, or by nucle-ophilic substitution by water or its reverse, and/or by an intramolecular molecular rearrangement.

This formal definition allows the original idea of oxidation (combination with oxygen), together with its extension to removal of hydrogen, as well as processes closely akin to this type of transformation, to be descriptively related to definition 1. For example, the oxidation of methane to chloromethane may be considered as follows:

CH4-2e- -h+ +oh- = Ch3oh^ reversal of hydrolysis CH3Cl.

Oxidation number is used to define the state of oxidation of an element. Unbound atoms have a zero oxidation state. The oxidation number of a central atom in a coordination entity is the charge it would bear if all the ligands were removed along with the electron pairs that were shared with the central atom (McNaught and Wilkinson, 1997). It is represented by a roman numeral, for example, Crm, CrVI.

The oxidation state is a measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-on set of rules: (l) the oxidation state of a free element (uncombined element) is zero; (2) for a simple (monoatomic) ion, the oxidation state is equal to the net charge on the ion; (3) hydrogen has an oxidation state of +1 and oxygen has an oxidation state of -2 when they are present in most compounds (exceptions to this are that hydrogen has an oxidation state of -1 in hydrides of active metals (e.g., LiH) and oxygen has an oxidation state of -1 in peroxides (e.g., H2O2); (4) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero. In ions, the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. For example, the oxidation states of sulfur in H2S, S8 (elementary sulfur), SO2, SO3, and H2SO4 are, respectively: -2, 0, +4, +6, and +6. The higher the oxidation state of a given atom, the greater is its degree of oxidation; the lower the oxidation state, the greater is its degree of reduction (McNaught and Wilkinson, 1997).

Another important reaction that metals undergo is the Fenton reaction, which is important in biological systems because most cells have some amounts of iron, copper, or other metals that can catalyze this reaction. Transition metals with redox potentials in a biologically accessible range, such as iron and copper, can accept and donate electrons in a catalytic fashion. The Fenton reactions results in generating oxidative species in the cell, as follows

This is the iron-salt-dependent decomposition of dihydrogen peroxide, generating the highly reactive hydroxyl radical, possibly by means of an oxoiron(IV) intermediate. Addition of a reducing agent, such as ascorbate, leads to a cycle that increases damage to biological molecules (McNaught and Wilkinson, 1997).

3.3 Inorganic Compounds

Metallic elements form a great variety of inorganic compounds. They can be classified into binary and multielement compounds. The most important binary compounds, both from the technological and the toxicological viewpoint, are oxides and sulfides. These are the chemical forms in which most metals appear in nature, the minerals and ores. The metal-con tain-ing aerosols produced in metallurgical processes often occur as metal oxides. However, in experimental toxi-cological studies, chlorides and acetates are the most commonly used metal compounds because of their high solubility in water and biological systems.

3.4 Metal Complexes

A metal complex or coordination compound is formed by the association of a metal atom or ion and another chemical species, called ligand, which may be either an anion or a polar molecule. The ligands such as BAL (2,3-dimercaptopropanol) and D-penicillamine ((CH3)2C(SH)CH(NH2)CO2H)) serve important biological functions, where -SH groups of the ligand easily bind to a metal. Because of this, they can be used as detoxifying agents in case of, for example, mercury exposure. Another example is EDTA (ethylenediami-netetraacetic acid) used in lead detoxification.

3.5 Organometallic Compounds

Organometallic compounds are classically compounds having bonds between one or more metal atoms and one or more carbon atoms of an organyl group. Organome-tallic compounds are classified by prefixing the metal with organo- (e.g., organopalladium compounds). In addition to the traditional metals and semimetals, elements such as boron, silicon, arsenic, and selenium are considered to form organometallic compounds. Examples are organomagnesium compounds MeMgl (iodo(methyl)magnesium); Et2Mg (diethylmagnesium); an organolithium compound BuLi (butyllithium); an organozinc compound ClZnCH2C( = O)O Et) chloro(eth oxycarbonylmethyl)zinc; an organocuprate Li+(CuMe2)-(lithium dimethylcuprate); and an organoborane Et3B (triethylborane). The status of compounds in which the canonical anion has a delocalized structure in which the negative charge is shared with an atom more electronegative than carbon, as in enolates, may vary with the nature of the anionic moiety, the metal ion, and, possibly, the medium. In the absence of direct structural evidence for a carbon-metal bond, such compounds are not considered to be organometallic (McNaught and Wilkinson, 1997).

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