Nucleus, atomic radius

Across a period, eff dominates. As we move across a period of main-group elements, electrons are added to the same outer level, so the shielding by inner electrons does not change. Because outer electrons shield each other poorly, Zeff on the outer electrons rises significantly, and so they are pulled closer to the nucleus. Atomic radius generally decreases in a period from left to right. 

Strictly speaking, the size of an atom is a rather nebulous concept The electron cloud surrounding the nucleus does not have a sharp boundary. However, a quantity called the atomic radius can be defined and measured, assuming a spherical atom. Ordinarily, the atomic radius is taken to be one half the distance of closest approach between atoms in an elemental substance (Figure 6.12). 

The decrease in atomic radius moving across the periodic table can be explained in a similar manner. Consider, for example, the third period, where electrons are being added to the third principal energy level. The added electrons should be relatively poor shields for each other because they are all at about the same distance from the nucleus. Only the ten core electrons in inner, filled levels (n = 1, n = 2) are expected to shield the outer electrons from the nucleus. This means that the charge felt by an outer electron, called the effective nuclear charge, should increase steadily with atomic number as we move across the period. As effective nuclear charge increases, the outermost electrons are pulled in more tightly, and atomic radius decreases. 

In order to explain his experimental results, Rutherford designed a new picture of the atom. He proposed that the atom occupies a spherical volume approximately I0 8cm in radius and at the center of each atom there is a nucleus whose radius is about 10 u cm. He further proposed that this nucleus contains most of the mass of the atom, and that it also has a positive charge that is some multiple of the charge on the electron. The region of space outside the nucleus must be occupied by the electrons. We see from Figure 14-11 that Rutherford s picture requires that most of the volume of the atom be a region of very low density. 

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. 

Atomic radii typically decrease from left to right across a period and increase down a group (Fig. 14.2 see also Fig. 1.46). As the nuclear charge experienced by the valence electrons increases across a period, the electrons are pulled closer to the nucleus, so decreasing the atomic radius. Down a group the valence electrons are farther and farther from the nucleus, which increases the atomic radius. Ionic radii follow similar periodic trends (see Fig. 1.48). 

The Lewis dot formalism shows any halogen in a molecule surrounded by three electron lone pairs. An unfortunate consequence of this perspective is that it is natural to assume that these electrons are equivalent and symmetrically distributed (i.e., that the iodine is sp3 hybridized). Even simple quantum mechanical calculations, however, show that this is not the case [148]. Consider the diiodine molecule in the gas phase (Fig. 3). There is a region directly opposite the I-I sigma bond where the nucleus is poorly shielded by the atoms electron cloud. Allen described this as polar flattening , where the effective atomic radius is shorter at this point than it is perpendicular to the I-I bond [149]. Politzer and coworkers simply call it a sigma hole [150,151]. This area of positive electrostatic potential also coincides with the LUMO of the molecule (Fig. 4). 

What was the importance of this research result for the chirality problem One difficulty is provided by the fact that the interaction responsible for the violation of parity is in fact not so weak at all, although it only acts across a very short distance (smaller than an atomic radius). Thus, the weak interaction is not noticeable outside the atomic nucleus, except for p-decay. It would thus have either no influence on chemical reactions or only a very limited effect on chemical reactions, as these almost completely involve only interactions between the electron shells. 

Consider the element with atomic number 116 in Group 6A. Even though it has not been isolated, its atomic radius is expected to be somewhat larger than that of Po (1.68 A), probably about 1.9 - 2.0 A, since it lies just below Po on the periodic table. Its outer electrons would lie in the n=l shell, which would be further away from the nucleus than Po s outermost electrons in the n=6 shell. 

Atomic radius The distance from the nucleus to the outer most electron orbital in an atom. 

One factor affecting atomic radii is changing n. As n increases, there is a higher prohahility of finding electrons farther from their nucleus. Therefore, the atomic volume is larger. In other words, atomic radius tends to increase with increasing n, and decrease with decreasing n. 

Ionization energy generally decreases down a group. Notice that this trend is the inverse of the trend for atomic radius. The two trends are, in fact, linked. As atomic radius increases, the distance of valence electrons from the nucleus also increases. There is a decrease, therefore, in the force of attraction exerted by the nucleus on the valence electrons. Thus, less energy is needed to remove one such electron. 

Ionization energy generally increases across a period. Again, this trend is linked to the atomic radius. Across a period, the atomic radius decreases because Zeff increases. The force of attraction between the nucleus and valence electrons is subsequently increased. Therefore, more energy is needed to remove one such electron. 

X 10-1 X cm. The atomic radius of atom is 1 A. If the mass number is 64, the ratio by the atomic volume that is occupied by the nucleus would be ... 

The radius of ail atom can be estimated by taking half the distance between the nucleus of two of the same atoms. For example, the distance between the nuclei of I2 is 2.66 A, half that distance would be the radius of atomic iodine or 1.33 A. Using this method the atomic radius ofneafcly all the elements can be estimated. 

FIGURE 5.20 Plots of atomic radius and Zeff for the highest-energy electron versus atomic number. As Zefr increases, the valence-shell electrons are attracted more strongly to the nucleus, and the atomic radius therefore decreases. 

The electronegativity of an atom depends on the radius of the atom. The atomic radius decreases and attraction exerted on valence electrons by the nucleus increases from left to right in a period. Atomic radius increases and attraction exerted on valence electrons by nucleus decreases from top to bottom. Therefore, electronegativity increases from left to right and decreases from top to bottom in the periodic table. 

The physical properties of the elements, such as melting point, boiling point and density are related to the atomic radius of the elements. Also, the atomic radius directly affects the ability of an atom to gain and lose electrons. The atomic radius is practically defined by assuming the shape of the atom as a sphere. The atomic radius is the distance between the nucleus and the outermost electron. But it is impossible to measure the atomic radius by separating the atoms from each other. Atomic Radius within a Group... 

Since the number of shells increases in the same group from top to bottom (by the period number increases), the atomic radius also increases. This means that the electron cloud around the nucleus becomes larger. The increase in the number of electrons causes them occupy a new energy level and orbitals. A higher energy level is always further from nucleus. Within a period, if the number of protons and electrons increases, the nuclear attraction force increases. This attraction force prevents an enormous increase in atomic radius. Atomic Radius Within a Period... 

As the atomic radius increases from top to bottom in a group, the valence electrons become more further away from the nucleus and the nuclear attraction forces on these electrons decrease. Therefore, as the atomic radius increases, the amount of energy required to remove an electron decreases. As a result, we can say that within a group ionization energy of elements decrease from top to bottom. 

The electron cloud around an atomic nucleus makes the concept of atomic size somewhat imprecise, but it is useful to refer to an atomic radius. One can arbitrarily divide the distance between centers of two bonded atoms to arrive at two radii, based on the crude picture that two bonded atoms are spheres in contact. If the bonding is covalent, the radius is called a covalent radius (see Table 8-2) if it is ionic, the radius is an ionic radius (see Table 9-2). The radius for non-bonded atoms may be defined in terms of the distance of closest non-bonding approach such a measure is called the van der Waals radius. These three concepts of size are illustrated in Figure 7-2. 

Within a given group of the periodic table, the first ionization energy decreases with increasing atomic number. This is related to the increase in atomic radius and the decreasing attraction of the nucleus for the increasingly distant outermost electron. It should be mentioned that this trend is not uniformly noted for the transition metals. 

The properties of the elements of the sixth period are influenced by lanthanide contraction a gradual decrease of the atomic radius with increasing atomic number from La to Lu. The elements of groups 5 to 11 for the fifth and sixth periods have comparable stmctural parameters. For instance, Nb and Ta, as well as the pair Mo and W, have very similar ionic radii, when they have the same oxidation number. As a result, it is very difficult to separate Nb and Ta, and it is also not easy to separate Mo and W. Similarly, Ag and Au have nearly the same atomic radius, 144 pm. Recent studies of the coordination compounds of Ag(I) and Au(I) indicate that the covalent radius of Au is even shorter than that of Ag by about 8 pm. In elementary textbooks the phenomenon of lanthanide contraction is attributed to incomplete shielding of the nucleus by the diffuse 4f inner subshell. Recent theoretical calculations conclude that lanthanide contraction is the result of both the shielding effect of the 4f electrons and relativistic effects, with the latter making about 30% contribution. 

The atomic radius is the distance between the nucleus and the outermost electron. Atomic radii are measured in nanometers (10 9 meters). In some fields, atomic radii are measured in a unit known as an angstrom, A (10 10 m or of a nanometer). Hydrogen is the smallest atom, measuring only 0.037 nm or 0.37 A. 

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