Atomic and physical properties of the elements

Some of the important properties of the elements are given in Table 18.1. The imprecision of the atomic weights of Kr and Xe reflects the natural occurrence of several isotopes of these elements. For He, however, and to a lesser extent Ar, a single isotope predominates ( He, 99.999 863% Ar, 99.600%) and much greater precision is possible. The natural preponderance of Ar is indeed responsible for the well-known inversion of atomic weight order of Ar and K in the periodic table, and the position of Ar in front of K was only finally accepted when it was shown that the atomic weight of He placed it in front of Li. The second isotope of helium, He, has only been available in significant amounts since 

CocKETT and K. C. Smith, Chap. 5 in Comprehensive Inorganic Chemistry, Vol. 1, pp. 139-211, Perga-mon Press, Oxford, 1973. G. A. Cook (ed.). Argon, Helium and the Rare Gases, 2 vols. Interscience, New York, 1961, 

Atomic and physical properties of the elements Table 18.1 Some properties of the noble gases 

All the elements have stable electronic configurations (Is or ns np ) and, under normal circumstances are colourless, odourless and tasteless monatomic gases. The non-polar, spherical nature of the atoms which this implies, leads to physical properties which vary regularly with atomic number. The only interatomic interactions are weak van der Waals forces. These increase in magnitude as the polarizabilities of the atoms increase and the ionization energies decrease, the effect of both factors therefore being to increase the interactions as the sizes of the atoms increase. This is shown most directly by the enthalpy of vaporization, which is a measure of the energy required to overcome the 

The stability of the electronic configuration is indicated by the fact that each element has the highest ionization energy in its period, though the value decreases down the group as a result of increasing size of the atoms. For the heavier elements is it actually smaller than for first-row elements such as O and F with consequences for the chemical reactivities of the noble gases which will be considered in the next section. Nuclear properties, particularly for xenon, have been exploited for nmr spectroscopy and Mdssbauer 

The Noble Gases Helium, Neon, Argon, Krypton, Xenon and Radon 

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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. 

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. 

For example, E. G. Mazurs (note 2, p. 105) expresses the discord as follows The periodicity of atomic structure must be accepted as a Natural Law. Therefore, scientists have to change their minds, get away from the conservatism that accepts only Mendeleev s chemical table as right, and adjust the other phenomena to this phenomenon that is, derive the chemical and physical properties of the elements from the electronic structure of the atoms. ... 

The periodic table was originally developed using atomic masses and chemical and physical properties of the elements, but it is now known that the atomic number is the real basis of the periodic table. 

Fundamental atomic and physical properties of the alkali metals are given in Tables 1, 2, and 3. The elements are characterized by having electron configurations each with a single s orbital electron outside a noble gas core (see Table 1). Sodium and cesium are mononucUdic so that their relative atomic masses are known extremely accurately in effect, the same can be written for potassium and rubidium since their isotopes (of which there are three and two, respectively) have... 

The number and arrangements of electrons in the outermost shells of atoms determine the chemical and physical properties of the elements as well as the kinds of chemical bonds they form. We write Lewis dot formulas (or Lewis dot representations, or just Lewis formulas) as a convenient bookkeeping method for keeping track of these chemically important electrons. We now introduce this method for atoms of elements in our discussion of chemical bonding in subsequent sections, we will frequently use such formulas for atoms, molecules, and ions. 

The Russian chemist, Dmitri Mendeleev, was a professor of chemistry at the University of St. Petersburg when he developed a periodic table of the elements. Mendeleev was studying the properties of the elements and realized that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements according to increasing atomic mass. For example, beryllium resembled magnesium, and boron resembled aluminum. Patterns of repeated properties began to appear. 

Krebs, Robert E. The History and Use of Our Earth s Chemical Elements A Reference Guide, 2nd ed. Westport, Conn. Greenwood Press, 2006. Following brief introductions to the history of chemistry and atomic structure, Krebs proceeds to discuss the chemical and physical properties of the elements group (column) by group. In addition, he describes the history of each element and current uses. 

Increasing the nuclear charge of an atom (together with its number of electrons) leads to the consecutive occupation by electrons of the electronic shells and subshells of higher and higher eneigy. This produces a quasi-periodicity (sometimes called periodicity in chemistry) of the valence shells, and as a consequence, a quasi-periodicity of all chemical and physical properties of the elements (reflected in the Mendeleev periodic table). 

The atomic number, the number of protons and hence the positive charge of the nucleus, determines the number of electrons that surround it. An electron has the same magnitude of electric charge as a proton, but opposite in sign. Therefore, for an atom to be electrically neutral the number of electrons outside the nucleus must be the same as the number of protons inside the nucleus. That is, the number of electrons is equal to the atomic number. Thus, hydrogen (atomic number 1) has one electron, carbon (atomic number 6) has six electrons, and so on, up to livermorium with its 116 electrons. Electrons are much lighter than protons and neutrons (by a factor of nearly 2,000), so their presence barely affects the mass of an atom. They have a profound effect on the chemical and physical properties of the element, and almost all chemistry can be traced to their behaviour. 

Atomic number (1821) n. The number (Z) of protons within the atomic nucleus. The electrical charge of these protons determines the number and arrangement of the outer electrons of the atom and thereby the chemical and physical properties of the element. 

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