The Periodic Table and Chemical Bonding

Three factors affect the properties of an atom or ion atomic number, mass number, and electron configuration. Process technicians are primarily concerned with the electron configuration and the valence shell. The outermost shell in an atom is referred to as the valence shell. The atoms of elements have different arrangements and they tend to react in an effort to fill the valence shell. The electrons in the valence shell are referred to as valence electrons. 

Dmitri Mendeleev, a popular Russian professor of chemistry who lived from 1834 to 1907, devised the first periodic table of elements. During his studies, Mendeleev recognized repetitions in the properties of elements that is, properties and similarities that repeated over and over again. This recurrence is referred to as periodicity or periodic. 

After some thought, Mendeleev numbered the columns with Roman numerals, some with the letter A and some with the letter B. This system was used virtually unchanged until quite recently, when minor variation between U.S. and European chemists (A/B variation) forced the international chemistry organization to recommend a column numbering system from 1-18, moving from left to right. In columns 1, 2, and 13-18, a number of elements extend above the rest. These elements are referred to as representative elements. The elements found in columns 3-12 are classified as transition elements. At the bottom of the periodic table are two rows that appear to be separated from the main body of the table. This is for convenience only. Elements 58-71 and 90-103 are called inner transition elements and actually fit between columns 3 and 4. 

The elements on the periodic table can be classified as metals, nonmetals, metalloids, or noble gases. Metals tend to be shiny and have atoms that give up electrons. Metals are malleable and tend to be excellent conductors of heat and electricity. By nature, nonmetals do not conduct electricity and have atoms that do not naturally give up electrons however, they do tend to accept electrons. The metalloid elements are located along the heavy black stair-step line on the right-hand side of the periodic table. Boron, silicon, germanium, arsenic, antimony, tellurium, polonium, and astatine are classified as metalloids. 

Covalent bonding is the mechanism of eiectron sharing that binds atoms together to form moie-cuies. in covaient bonding, a bond is made of a pair of eiectrons shared by two atoms. More com-piex eiectron structures may share one, two, or three eiectron pairs between atoms. Exampies of this inciude  

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In earlier courses, the basics of atomic structure, the periodic table, and chemical bonding are investigated. The first two columns of Figure 2.1 are a chronological display of some of the concepts often discussed. 

Carbon has by far the greatest number of compounds of any element. The thousands of combinations of carbon with other elements give it the diversity of compounds that makes it the basis of life. Carbon has a reacting power of 4 as it has four electrons in its outermost shell and is placed in group 4 of the periodic table. Its chemical bonds are all covalent. There are over half a million compounds of carbon and hydrogen alone and some are very useful, including the hydrocarbons in petrol. 

The concept of an octet of electrons is one of the foundations of chemical bonding. In fact, C, N, and O, the three elements that occur most frequently in organic and biological molecules, rarely stray from the pattern of octets. Nevertheless, an octet of electrons does not guarantee that an inner atom is in its most stable configuration. In particular, elements that occupy the third and higher rows of the periodic table and have more than four valence electrons may be most stable with more than an octet of electrons. Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons. In the third row, phosphoms, with five valence electrons, can form as many as five bonds. Sulfur, with six valence electrons, can form six bonds, and chlorine, with seven valence electrons, can form as many as seven bonds. 

Equation 16.12 expresses a relation between q and B.This is not a universal relation, but it does apply to the sp-bonded elements of the first four columns of the Periodic Table. Using chemical hardness values given by Parr and Yang (1989), and atomic volumes from Kittel (1996), it has been shown that the bulk moduli of the Group I, II, III, and IV elements are proportional to the chemical hardness density (CH/atomic volume) (Gilman, 1997). The correlation lines pass nearly through the coordinate origins with correlation coefficients, r = 0.999. Thus physical hardness is proportional to chemical hardness (Pearson, 2004). 

Quantum chemists have developed considerable experience over the years in inventing new molecules by quantum chemical methods, which in some cases have been subsequently characterized by experimentalists (see, for example, Refs. 3 and 4). The general philosophy is to explore the Periodic Table and to attempt to understand the analogies between the behavior of different elements. It is known that for first row atoms chemical bonding usually follows the octet rule. In transition metals, this rule is replaced by the 18-electron rule. Upon going to lanthanides and actinides, the valence f shells are expected to play a role. In lanthanide chemistry, the 4f shell is contracted and usually does not directly participate in the chemical bonding. In actinide chemistry, on the other hand, the 5f shell is more diffuse and participates actively in the bonding. 

Electronegativity is greatest for elements at the upper right of the periodic table and lowest for elements at the lower left. Noble gases are not considered in electronegativity discussions because, with only a few exceptions, they do not participate in chemical bonding. 

One feature of metals is well known. Metals tend to lose electrons to nonmetals in a chemical reaction. That is, they tend to have lower electronegativities than nonmetals. This is obvious in compounds formed from metals at the far left of the periodic table and nonmetals from the far right. Sodium (a metal) clearly loses an electron to chlorine (a nonmetal) forming an ionic bond. The resulting compound—table salt—is a water-soluble, white... 

Selenium is in group 16 of the periodic table and although it has chemical and physical properties intermediate between metals and nonmetals (Table 1), it is usually described as a nonmetal. The chemical behavior of selenium has some similarities to that of sulfur. Formally, selenium can exist in the —II, 0, IV, and VI oxidation states (Table 2). Selenium has six natural stable isotopes, the most important being Se and °Se. Although Se is generally also regarded as a stable isotope, it is a )8-emitter with a very long half-life (1.4 X 10 ° yr). Both arsenic and selenium tend to be covalently bonded in all of their compounds. 

Transition metal carbides and phosphides have shown potential as highly active catalysts. In these compounds, the C and P sites cannot be considered as simple spectators. They moderate the reactivity of the metal centers and provide bonding sites for adsorbates. The reactivity of the C centers in MC(OOl) surfaces varies in a complex way with the position of the metal in the Periodic Table and the filling of the carbide valence band. M Cj metcars should display a catalytic performance even better than that of the well-known Mo C or MC catalysts. By introducing six pairs of groups in the structure, the system is stabilized, while the presence of four low-coordinated M sites allows a reasonably high chemical reactivity. 

The subshells (see Fig. 5.12) determine the structure of the periodic table and the formation of chemical bonds. In preparation for a discussion of these connections, it is necessary to describe the energy values for Hartree orbitals. 

The Hartree orbitals are the foundation of the quantum explanation of atomic structure. They justify the shell model of the atom, they explain the structure of the periodic table, and they provide the starting point for the quantum explanation of chemical bond formation in the following chapter. 

The chemical formula NaCl does not explicitly indicate the ionic nature of the compound, only the ratio of ions. Furthermore, values of electronegativities are not always available. So we must learn to recognize, from positions of elements in the periodic table and known trends in electronegativity, when the difference in electronegativity is large enough to favor ionic bonding. 

Inorganic chemistry is concerned with the chemical elements (of which there are about 100) and the extremely varied compounds they form. The essentially descriptive subject matter is unified by some general concepts of structure, bonding and reactivity, and most especially by the periodic table and its underlying basis in atomic structure. As with other books in the Instant Notes series, the present account is intended to provide a concise summary of the core material that might be covered in the first and second years of a degree-level course. The division into short independent topics should make it easy for students and teachers to select the material they require for their particular course. 

In 1902, while explaining the laws of valence to his students at Harvard, Lewis conceived a concrete model for this process, something Abegg had not done. He proposed that atoms were composed of a concentric series of cubes with electrons at each of the resulting eight comers. This cubic atom explained the cycle of eight elements in the Periodic Table and corresponded to the idea that chemical bonds were formed by the transfer of electrons so each atom had a complete set of eight electrons. Lewis did not publish his theory, but fourteen years later it became an important part of his theory on the shared electron-pair bond. 

Both surface atoms and adsorbates must participate to form the surface chemical bond. In order to determine the nature of the bond, the heat of adsorption is measured as a function of the pertinent variables. These include trends across the periodic table, variations of bond energies with adsorbate size, molecular structure and coverage, and substrate structure. Changes in the electronic and atomic structure of the bonding partners are determined and compared with their electronic and atomic (or molecular) structure before they formed the surface bond. 

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