Shells second electron

If certain quanta suitable for the excitation of a line are absorbed without photon emission, a radiationless transition is likely. This transition is known as the Auger effect,39 and it may be thought to involve an absorption by the atom of the photon produced when the hole in the K shell is filled by an electron from one of the external shells such as the L shell. The absorption of this photon results in the ejection of a second electron from one of the shells to leave a doubly charged residue of what had been a normal atom. The atom in this condition is described by naming the two states in which the electron holes are to be found e.g., the atom is in the LL or LM or LN state. An atom in such a state is, of course, vastly different from the usual divalent cation. 

Redress can be obtained by the electron localization function (ELF). It decomposes the electron density spatially into regions that correspond to the notion of electron pairs, and its results are compatible with the valence shell electron-pair repulsion theory. An electron has a certain electron density p, (x, y, z) at a site x, y, z this can be calculated with quantum mechanics. Take a small, spherical volume element AV around this site. The product nY(x, y, z) = p, (x, y, z)AV corresponds to the number of electrons in this volume element. For a given number of electrons the size of the sphere AV adapts itself to the electron density. For this given number of electrons one can calculate the probability w(x, y, z) of finding a second electron with the same spin within this very volume element. According to the Pauli principle this electron must belong to another electron pair. The electron localization function is defined with the aid of this probability ... 

Suppose we want to write the electron configuration of scandium (atomic number 21). We can rewrite the first 12 electrons that we wrote above for magnesium, and then just keep going. As we added electrons, we filled the first shell of electrons first, then the second shell. When we are filling the third shell, we have to ask if the electrons with n = 3 and / = 2 will enter before the n = 4 and 1 = 0 electrons. Since (n + /) for the former is 5 and that for the latter is 4, we must add the two electrons with n = 4 and / = 0 before the last 10 electrons with n = 3 and / = 2. In this discussion, the values of m and s tell us how many electrons can have the same set of n and / values, but do not matter as to which come first. 

Because of their having larger sizes and more filled shells of electrons between the outer shell and the nucleus, the ionization energies of second- and third-row metals are lower than those of first-row metals. Consequently, it is easier for the heavier metals to achieve higher oxidation states, which also favors higher coordination numbers. In general, there is also a greater tendency of the heavier metals... 

The nomenclature for the emitted x-rays is related to the shells involved in the transition, as illustrated in Fig. 10b. For example, if the initial electron is ejected from the K shell and the second electron drops from the adjacent L shell, a K-a x-ray is emitted. If the electron drops from the M shell, the emitted x-ray is a K-fi x-ray. The most common lines observed are the KIM lines. 

The sixth element in the periodic table, carbon, has the electron configuration 2s 2 and, thus, has 4 valence electrons in the unfilled orbitals of its second electron shell. To fill these orbitals to a stable set of 8 valence electrons, a single carbon atom may share electrons with 2, 3, or even 4 other atoms. No other element forms such strong bonds to as many other atoms as carbon does. Moreover, multiple carbon atoms readily link together with single, double, or triple bonds. These factors make element number 6 unique in the entire periodic table. The number of carbon-based compounds is many times greater than the total of all compounds lacking carbon. 

To account for this behaviour these authors proposed an extension of the theory, given by Pople and Santry, which predicted that the coupling constant depends upon the energies of the s relative to those of the p, d, etc. electrons of the coupled nuclei and additionally upon the resonance integral between the outer-shell r electrons of the two atoms. It was considered that the first factor remained fairly constant for the P-P bonds and the observed variation was qualitatively explained by the dependence of the second factor upon the electronegativity and the bulkiness of the substituent. 

Beyond the first shell, the electrons are further divided among sub-shells. The second shell has one sub-shell of two electrons and one of six. The third shell has one sub-shell of two, one of six, and one of ten. The fourth shell has sub-shells containing 2, 8,10, and 14 electrons. 

Carefully consider the differences between these two compounds. Where the potassium fluoride contains a potassium ion, the molecular fluorine contains a fluorine. What then arc the differences between the potassium ion and the fluorine atom First consider size. The potassium ion has three full shells of electrons, which makes it larger. The distance between the two nuclei of the fluorine atoms within F, therefore, should be closer together. Second, consider the type of bonding. The F, compound is covalent. This involves the overlapping of shells, which allows the nuclei to be closer. So, by both considerations, the nuclei of molecular fluorine, F2, should be closer together. 

In this reaction (demonstrated in vitro), one of the two radicals is oxidizing while the other is reducing. In vivo, this reaction is catalyzed by one of several isoforms of an enzyme known as superoxide dismutase (SOD). As shown above, hydrogen peroxide may form as a result of the superoxide anion s dismutation reaction however, it may also be produced from a bivalent reduction of 02. The addition of the second electron leads to the formation of hydrogen peroxide, which is a powerful oxidizing agent. Due to the unpaired electrons in their outer shells, free radicals are favored to pair with other molecules during bimolecular collisions. 

The atomic radii of the second- and third-series transition elements from group 4B on are nearly identical, though we would expect an increase in size on adding an entire principal quantum shell of electrons. The small sizes of the third-series atoms are associated with what is called the lanthanide contraction, the general decrease in atomic radii of the /-block lanthanide elements between the second and third transition series (Figure 20.4). 

The lanthanide contraction is due to the increase in effective nuclear charge with increasing atomic number as the 4/ subshell is filled. By the end of the lanthanides, the size decrease due to a larger Zeff almost exactly compensates for the expected size increase due to an added quantum shell of electrons. Consequently, atoms of the third transition series have radii very similar to those of the second transition series. 

There is one more complication to the electron shells. Inside the shells themselves, electrons can be found in regions called orbitals. There are four types of orbitals—s, p, d, and/—and each has a specific shape. Blocks of the periodic table correspond to the different orbitals. The electrons in atoms of the first row of the table are found in the Is orbital. Helium, at the far right of the first row, consists of 2 electrons in the Is orbital. Neon, at the far right of the second row, has two electrons in the Is orbital, 2 electrons in the 2s orbital, and 6 electrons in the 2p orbital. These arrangements of electrons within orbitals are known as electron configurations. Chemists notate the electron configuration of helium as Is2 and neon as ls22s22p6. 

Each atom in both structures has a complete valence shell of electrons. There are no formal charges in the first structure, but in the second structure, the oxygen is formally positive and the carbon is formally negative. 

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