Isoelectronic ions, sizes

Sizes of ions related to positions of elements in the periodic table. Note that size generally increases down a group. Also note that in a series of isoelectronic ions, size decreases with increasing atomic number. The ionic radii are given in units of picometers. 

For isoelectronic ions, size generally decreases as 1 increases. 

Much of the chemistry of oxygen can be rationalized in terms of its electronic structure (2s 2p ), high electronegativity (3.5) and small size. Thus, oxygen shows many similarities to nitrogen (p. 412) in its covalent chemistry, and its propensity to form H bonds (p. 52) and p double bonds (p. 416), though the anionic chemistry of 0 and OH is much more extensive than for the isoelectronic ions N , and NH2. Simi-... 

Table 3.3 gives a compilation of ionic radii. It is seen that for isoelectronic ions the radius decreases as the net positive charge increases. In relation to the periodic table this means that ion size decreases from left to right in a period. This is due to the same factors that make the ionization potential increase electrons are held more tightly and thus move closer to the nucleus. Going down a periodic family, we find that the increasing principal quantum number and the increased screening by core electrons produce an increase in ionic size. 

It is instructive to consider the Stokes shift of the emission. It depends on the coupling parameter S like the occurrence of vibrational structure does. Table 3 shows for three different 5s2 ions, Sn2+, Sb3+ and Te4+, in comparable host lattices the Stokes shift and the value of the spin-orbit coupling constant [19]. The luminescent ions have practically the same size as the host lattice ions involved. It is clear from the table that the Stokes shift increases drastically in the sequence of isoelectronic ions Te4+ [Pg.13]

One trend worth noting involves the relative sizes of a set of isoelectronic ions—ions containing the same number of electrons. Consider the ions O2-, F, Na+, Mg2+, and Al3+. Each of these ions has the neon electron configuration. How do the sizes of these ions vary In general, there are two important facts to consider in predicting the relative sizes of ions the number of electrons and the number of protons. Since these ions are isoelectronic, the number of electrons is 10 in each case. Electron repulsions should therefore be about the same in all cases. However, the number of protons increases from... 

The isoelectronic ions C104, SO4 ", and P04 have been studied by Johansen. " A medium size basis set was used, and regular tetrahedral geometry assumed. The energy-level order was in agreement with experiment. Other work on these molecules particularly concerned with the photoelectron spectrum has been reported. ... 

Decreasing size with increasing nuclear charge for isoelectronic ions... 

Table 4.6 The sizes of a series of isoelectronic ions, with the [Ne] inert gas core...

Figure 8.9 shows the radii of ions derived from the familiar elements, arranged according to elements positions in the periodic table. We can see parallel trends between atomic radii and ionic radii. For example, from top to bottom both the atomic radius and the ionic radius increase within a group. For ions derived from elements in different groups, a size comparison is meaningful only if the ions are isoelectronic. If we examine isoelectronic ions, we find that cations are smaller than anions. For example, Na is smaller than F . Both ions have the same number of electrons, but Na... 

One trend worth noting involves the relative sizes of a set of isoelectronic ions—ions containing the same number of electrons. Consider the ions F, Na", Mg, and AP. ... 

There are similar challenges in determining the size of ions. Because the stable ions of the different elements have different charges and different numbers of electrons, as well as different crystal structures for their compounds, it is difficult to find a suitable set of numbers for comparison. Earlier data were based on Pauling s approach, in which the ratio of the radii of isoelectronic ions was assumed to be equal to the ratio of their effeetive... 

For isoelectronic ions, the size of the ion is based on the size of the electron cloud, not on the number of protons in the nucleus. 

Despite this drawback there is an interest in science and technology in single-ion quantities for rationalizing the discussion of electrolyte solution properties. Various methods have been developed for their estimation based on extrathermodynamic assumptions, such as the following (1) The contributions of cation and anion are set equal for a salt composed of ions of equal charge and approximately equal radii (2) the results of measurements on a series of homologous electrolytes are extrapolated with regard to ionic radii or ionic volumes to zero ion size or zero reciprocal radii (3) the differences in conventional ionic properties are used for theoretically rationalized extrapolations and (4) the properties of ions are compared with those of isoelectronic neutral molecules of similar chemical constitution and size. 

Focusing on isoelectronic cations, the radii of tripositive ions (ions that bear three positive charges) are smaller than those of dipositive ions (ions that bear two positive charges), which in turn are smaller than unipositive ions (ions that bear one positive charge). This trend is nicely illnstrated by the sizes of three isoelectronic ions in the third period Al has... 

The charges of a set of isoelectronic ions vary from 3-h to 3 —. Place the ions in order of increasing size. 

The fluoride ion is the least polarizable anion. It is small, having a diameter of 0.136 nm, 0.045 nm smaller than the chloride ion. The isoelectronic E and ions are the only anions of comparable size to many cations. These anions are about the same size as K" and Ba " and smaller than Rb" and Cs". The small size of E allows for high coordination numbers and leads to different crystal forms and solubiUties, and higher bond energies than are evidenced by the other haUdes. Bonds between fluorine and other elements are strong whereas the fluorine—fluorine bond is much weaker, 158.8 kj/mol (37.95 kcal/mol), than the chlorine—chlorine bond which is 242.58 kJ/mol (57.98 kcal/mol). This bond weakness relative to the second-row elements is also seen ia 0-0 and N—N single bonds and results from electronic repulsion. 

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