Atomic radius/radii nuclear atom

Electronegativity is the tendency of an atom to attract the bonding electrons within a compound to itself. It depends upon the nuclear charge (proton number) and the atomic radius of the atom. It is these factors that control the ionization energy of the atom which in turn is related to the ability of an atom to attract electrons. 

The correct answer is (A). Fluorine has the smallest atomic radius. The fluorine atom has the highest effective nuclear charge of the elements in the list. Because there are no elements on the list with a greater effective nuclear charge or a smaller amount of shielding, fluorine will have the smallest atomic radius. 

The Electron Affinity for an element depends on the atomic radius, the nuclear charge, and the screening effect of inner layers of electrons. 

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. 

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 atomic radii of the second row of d-metals (Period 5) are typically greater than those in the first row (Period 4). The atomic radii in the third row (Period 6), however, are about the same as those in the second row and smaller than expected. This effect is due to the lanthanide contraction, the decrease in radius along the first row of the / block (Fig. 16.4). This decrease is due to the increasing nuclear charge along the period coupled with the poor shielding ability of /-electrons. When the d block resumes (at lutetium), the atomic radius has fallen from 217 pm for barium to 173 pm for lutetium. 

As we move across the period, atomic radius decreases as electrons are added to the same shell ( = 3) (1) and, therefore, the outermost electrons experience the same shielding (1) from the nuclear charge. The increase in the number of protons (1) causes the decrease in radius from Na to Ar. These factors also explain the general increase in first ionisation energy across the period. 

These three structures are the predominant structures of metals, the exceptions being found mainly in such heavy metals as plutonium. Table 6.1 shows the structure in a sequence of the Periodic Groups, and gives a value of the distance of closest approach of two atoms in the metal. This latter may be viewed as representing the atomic size if the atoms are treated as hard spheres. Alternatively it may be treated as an inter-nuclear distance which is determined by the electronic structure of the metal atoms. In the free-electron model of metals, the structure is described as an ordered array of metallic ions immersed in a continuum of free or unbound electrons. A comparison of the ionic radius with the inter-nuclear distance shows that some metals, such as the alkali metals are empty i.e. the ions are small compared with the hard sphere model, while some such as copper are full with the ionic radius being close to the inter-nuclear distance in the metal. A consideration of ionic radii will be made later in the ionic structures of oxides. 

B. It is significantly more difficult to remove neon s most loosely held electron (Ii) than that of beryllium s I,. This trend is also noted when examining I2 s and I3 s. Neon also has a greater nuclear charge than beryllium, which, if all factors are held constant, would result in a smaller atomic radius. 

A) The atomic radius decreases because of increasing effective nuclear charge and electrostatic attraction. There are more protons and electrons, so electrons are needed to create a complete shell thus, there is an increase in electronegativity. 

How really means how and why. Questions that ask how one variable is affected by another—and these questions are legion—require an explanation, even if the question doesn t seem to specifically ask how and why. For example, you might be asked to explain how effective nuclear charge affects the atomic radius. If you say that the atomic radius decreases, you may have only received one of two possible points. If you say that this is because effective nuclear charge has increased, you can earn the second point. 

According to Gordy (1946), electronegativity is represented by the value of the potential resulting from the effect of the nuclear charge of an unshielded atom on a valence electron located at a distance corresponding to the covalent radius of the atom. 

The covalent radius between identical atoms also decreases within a period when the group number is increased, due to the larger nuclear charges exerting more attraction on the electrons (table 4.10). 

I Neon Neon also has a greater nuclear charge than beryllium, which if all factors are held constant, would result in a smaller atomic radius. 

The values of ioni/ulion energies and atomic sizes are influenced by retain islic dlccls that, for valence electrons, increase with the value of 1 /. and become sufficiently important in the elements of the 6lh period (C s Rn) to explain largely their chemical differences from the elements of the 5lli period (Rh- Xe). The initial relativistic effect is to cause a decrease in the radius of the 1 s atomic orbital of Ihe atom. The I mass of the electron in the Is orbital becomes higher as the nuclear charge increases because the velocity of the electron increases. 

The hydrogen atom has a nuclear charge of unity and therefore has one electron. According to Bohr this electron will have one velocity so that it moves in a circular path with a radius of... 

Nuclear Atom. From the results of the experiments, Rutherford concluded that the mass in the positive charge of an atom, instead of being distributed throughout the volume of a sphere of the order of 10-3 centimeter in radius, was concentrated in a very small volume of the order of 10-12 centimeter in radius, He thus developed the idea of a nuclear atom. I he atom was pictured as a small solar system with the very heavy and highly charged nuclens occupying the position of the sun, and with electrons moving around it, as planets in their respective orbits. 

Ionization energies generally increase from left to right across a transition series, though there are some irregularities, as indicated in Table 20.1 for the atoms of the first transition series. The general trend correlates with an increase in effective nuclear charge and a decrease in atomic radius. 

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... 

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