Trends in atomic and physical properties

In so far as the chemical (and physical) properties of an element derive from its electronic configuration, and especially the configuration of its least tightly bound electrons, it follows that chemical periodicity and the form of the periodic table can be elegantly interpreted in terms of electronic structure. 

General similarities and trends in the chemical properties of the elements had been noticed increasingly since the end of the eighteenth century and predated the observation of periodic variations in physical properties which were not noted until about 1868. However, it is more convenient to invert this order and to look at trends in atomic and physical properties first. 

Chapter 30). Similar plots are obtained for the atomic and ionic radii of the elements and an inverted diagram is obtained, as expected, for the densities of the elements in the solid state (Fig. 2.2). 

Of more fundamental importance is the plot of first-stage ionization energies of the elements, i.e. the energy /m required to remove the least tightly bound electron from the neutral atom in the gas phase  

These are shown in Fig. 2.3 and illustrate most convincingly the various quantum shells and subshells described in the preceding section. The energy required to remove the I electron from an atom of hydrogen is 13.606 eV (i.e. 1312 kJ per mole of H atoms). This rises to 2372 kJ mol for He (Is-) since the positive charge on the helium nucleus is twice that of the 

Of more fundamental importance is the plot of first-stage ionization energies of the elements, 

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These, though more difficult to describe quantitatively than the trends in atomic and physical properties described in the preceding subsection, also become apparent when the elements are compared in each group and along each period. Such trends will be discussed in detail in later chapters and it is only necessary here to enumerate briefly the various types of behaviour that frequently recur. 

The trends in chemical and physical properties of the elements described beautifully in the periodic table and the ability of early spectroscopists to fit atomic line spectra by simple mathematical formulas and to interpret atomic electronic states in terms of empirical quantum numbers provide compelling evidence that some relatively simple framework must exist for understanding the electronic structures of all atoms. The great predictive power of the concept of atomic valence further suggests that molecular electronic structure should be understandable in terms of those of the constituent atoms. 

The periodic table has been described as the chemist s best friend. Chemical reactions involve loss, gain, or sharing of electrons. In this chapter, we have seen that the fundamental basis of the periodic table is that it reflects similarities and trends in electron configurations. It is easy to use the periodic table to determine many important aspects of electron configurations of atoms. Practice until you can use the periodic table with confidence to answer many questions about electron configurations. As we continue our study, we will learn many other useful ways to interpret the periodic table. We should always keep in mind that the many trends in chemical and physical properties that we correlate with the periodic table are ultimately based on the trends in electron configurations. 

Some of the most common elements foimd on Earth occupy positions in the first two columns of the periodic table. The elements in the first column, the alkali metals, exhibit regular trends in chemical and physical properties that are apparent as the size of their atoms increases. Elements in the second column, the alkaline earths, also exhibit similar... 

Several atomic and physical properties of the elements are given in Table 16.2. The trends to larger size, lower ionization energy and lower electronegativity are as expected. The trend to metallic conductivity is also noteworthy indeed, Po resembles its horizontal neighbours Bi, Pb and T1 not only in this but in its moderately high density and notably low mp and bp. 

Dalton s little hard sphere model of the atom may seem primitive by today s standards, but it was an essential step in the evolution of chemical knowledge. Dalton s model persisted for almost one hundred years before anyone could think of any way to improve upon it. What is especially remarkable is that Dalton s theory was not completely accepted by the scientific community. Until 1900, there remained prominent physicists and chemists who continued to deny the existence of atoms. Actually, probably the most unsatisfying thing about Dalton s model is that it offered no explanation for the differences in chemical and physical properties that were observed among the elements. Even Dmitri Mendeleev, who, in 1869, developed the modern periodic table of the elements, could offer no explanation for the regular, or periodic, trends in the elements that were displayed in his periodic table. For that explanation, we must turn the clock forward to the events of the 1890s. 

All the elements in a main group have in common a characteristic valence electron configuration. The electron configuration controls the valence of the element (the number of bonds that it can form) and affects its chemical and physical properties. Five atomic properties are principally responsible for the characteristic properties of each element atomic radius, ionization energy, electron affinity, electronegativity, and polarizability. All five properties are related to trends in the effective nuclear charge experienced by the valence electrons and their distance from the nucleus. 

A comparison of steric influence resulting from more than 80 different silyl groups is summarized in Table 1. The trend is established on the basis of reaction yields, rates and selectivity, as well as on reactivity and physical properties of organosilanes. Two different effects may be involved (a) the effect of a silyl group on reactions taking place at the neighboring centers, and (b) the effect of the groups attached to silicon on the nucleophilic attack at the silicon atom. 

There are two main classes of isomers. Figure 22-10 shows compounds that are examples of structural isomers. The atoms of structural isomers are bonded in different orders. The members of a group of structural isomers have different chemical and physical properties despite having the same formula. This observation supports one of the main principles of chemistry The structure of a substance determines its properties. How does the trend in boiling points of C5H12 isomers relate to their molecular structures ... 

Table 8.1 lists some key physical properties of the elements of the first transition series, taken mostly from Appendix F. The general trends in all of these properties can be understood by recalling that nuclear charge also increases across a period as electrons are being added to the same subshell, in this case, the d shell. The first and second ionization energies tend to increase across the period, but not smoothly. The energies of the 4s and 3d orbitals are so close to one another that the electron configurations of the neutral atoms and their ions are not easily predicted from the simplest model of atomic structure. 

After Dalton, there were several attempts throughout Western Europe to organize the known elements into a conceptual framework that would account for the similar properties that related groups of elements exhibit and for trends in properties that correlate with increases in atomic weights. The most successful periodic table of the elements was designed in 1869 by a Russian chemist, Dmitri Mendeleev. Mendeleev s method of organizing the elements into columns grouping elements with similar chemical and physical properties proved to be so practical that his table is still essentially the only one in use today. 

When the Russian chemist Dmitri Mendeleev (1834-1907) developed his version of the periodic table in 1869, he arranged the elements known at that time in order of atomic mass or atomic weight so that they fell into columns called groups or families consisting of elements with similar chemical and physical properties. By doing so, the rows exhibit periodic trends in properties going from left to right across the table, hence the reference to rows as periods and name periodic table. ... 

By building on the work of Mendeleev and others, and by using the concept of the atomic number, we are now able to state the modern periodic law When elements are arranged in the order of their atomic numbers, thdr chemical and physical properties show repeatable, or periodic, trends. Other familiar periodic phenomena include the changing of the seasons and the orbits of the planets, which are periodic with time. To illustrate this idea on a simple level, a shingle roof, which has the same pattern over and over, is periodic. 

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