P Subshell

An effective way to determine the detailed electron configuration of any element is to use the periodic table to determine which subshell to fill next. Each s subshell holds a maximum of 2 electrons each p subshell holds a maximum of 6 electrons each d subshell holds a maximum of 10 electrons and each / subshell holds a maximum of 14 electrons (Table 17-5). These numbers match the numbers of elements in a given period in the various blocks. To get the electron configuration, start at hydrogen (atomic number = 1) and continue in order of atomic number, using the periodic table of Fig. 17-10. 

Ans. A state of great stability is a state in which the outermost s and p subshells are filled and no other subshell of the outermost shell has any electrons. 

The order of subshell filling is s2s2p3s3p4s3d4p5s4d5p6s4f 5d6pls5f 6d. An s subshell can have a maximum of 2 electrons, a p subshell can have 6 electrons maximum, a d subshell can have 10, and an / subshell can have 14. 

There are several ways of indicating the arrangement of the electrons in an atom. The most common way is the electron configuration. The electron configuration requires the use of the n and / quantum numbers along with the number of electrons. The principle quantum number, n, is represented by an integer (1,2,3. ..), and a letter represents the l quantum number (0 = s, 1 = p, 2 = d, and 3 = f). Any s-subshell can hold a maximum of two electrons, any p-subshell can hold up to six electrons, any d-subshell can hold a maximum of 10 electrons, and any f-subshell can hold up to 14 electrons. 

Maximum number of electrons for s-subshells = 2, p-subshells = 6, d-subshells = 10, f-subshells = 14. 

As the first shell has only two electrons, it has only one orbital - the Is orbital. The second and subsequent shells all have p orbitals in addition to an s orbital. The p subshell has three different p orbitals of the same energy. Orbitals that have the same energy are said to be degenerate. 

Electrons fill the orbitals in order of increasing energy, meaning that the lowest energy subshells are filled first. This is known as the aufbau principle. Of course, some subshells, such as the p subshell and the d subshell, have degenerate orbitals. 

The diagram above shows the electronic configuration for carbon in orbital box notation. The two electrons in the p subshell are in different orbitals, but have parallel spins, and the electrons sharing the same orbitals in the Is and 2s subshells have opposite spins. The diagram also suggests that one of the 2p orbitals is empty. In reality, there is no such thing as an empty orbital. If an orbital is empty, then it does not exist. However, it is acceptable to show empty orbitals in this type of notation. 

This can be explained in terms of the relative stability of different electronic configurations and thus provides evidence for these electronic configurations. To help you understand this, you have to appreciate that there is a special stability associated with a filled subshell or a half-filled subshell - for example, the p subshell when it contains three or six electrons. Likewise, the d subshell is most stable when it contains five or ten electrons. The more stable the electronic configuration, then the more difficult it is to remove an electron and therefore the ionisation energy is higher. 

There will be two electrons in the s subshell, six electrons in the p subshell and 10 electrons in the d subshell, which means that there must be 14 electrons in the f subshell. Therefore there must be seven f orbitals to accommodate these 14 electrons. 

The stabilization of a half-filled d shell is even more pronounced than tltat of the p subshell. Why ... 

Orbitals can be grouped into successive layers, or shells, according to their principal quantum number n. Within a shell, orbitals are grouped into s, p, d, and f subshells according to their angular-momentum quantum numbers l. An orbital in an s subshell is spherical, an orbital in a p subshell is dumbbell-shaped, and four of the five orbitals in a d subshell are cloverleaf-shaped. 

The expansion that occurs when a group 7A atom gains an electron to yield an anion can t be accounted for by a change in the quantum number of the valence shell, because the added electron simply completes an already occupied p subshell [Ne] 3s2 3p5 for a Cl atom becomes [Ne] 3s2 3p6 for a Cl- anion, for example. Thus, the expansion is due entirely to the decrease in effective nuclear charge and the increase in electron-electron repulsions that occurs when an extra electron is added. 

In looking for other trends in the data of Figure 6.6, the near-zero Eea s of the alkaline earth metals (Be, Mg, Ca, Sr, Ba) are particularly striking. Atoms of these elements have filled s subshells, which means that the additional electron must go into a p subshell. The higher energy of the p subshell, together with a relatively low Zeff for elements on the left side of the periodic table, means that alkaline earth atoms accept an electron reluctantly and have Eea values near zero. 

The tendency of main-group atoms to fill their s and p subshells when they form bonds—the octet rule discussed in Section 6.12—is an important guiding principle that makes it possible to predict the formulas and electron-dot structures of a great many molecules. As a general rule, an atom shares as many of its valence-shell electrons as possible, either until it has no more to share or until it reaches an octet configuration. For second-row elements in particular, the following guidelines apply ... 

The electrons in an uncharged arsenic atom (As0) are located in the s subshell of the first principal quantum number (n = 1), the s and p subshells of principal quantum numbers 2-4 (n = 2-4), and the d subshell of the third principal quantum number (n = 3). Specifically, the As0 electron configuration may be written as ... 

The s subshells have one orbital, the p subshells have three, and the d subshell has five. Each orbital may contain up to two electrons. For example, the 2p subshell has a total of six electrons, where each of the three 2p orbitals contains two electrons (Faure, 1998), 63-71. 

If we examine the electrons of a lone carbon atom in its ground state we would see that its four valence electrons are in their expected atomic orbitals, two in the orbital of the subshell and two in orbitals of the p subshell. The p electrons are at a higher energy state than the electrons. 

All orbitals with a value of 1 = 1, are the orbitals of the p subshell. The p orbitals are not spherical. Each p subshell consists of three orbitals in the form of lobes that differ in their orientation. These lobes are separated from each other by a plane where the probability of finding the electron is zero. The lobes are located on both sides of this plane like a dumbbell. The shape of the p orbitals are the same but the directions of the lobes are different. Since it is possible to imagine that these lobes are oriented along x, y and z coordinates, so the corresponding p orbitals are denoted by px, py and pz. Hence, in all main energy levels, except first energy level (n = 1), there are three p orbitals. 

Some free atoms can capture an extra electron to form a stable gaseous anion, particularly with elements having almost-completed p subshells (Groups VIA and VIIA, especially). For example,... 

We have seen that the maximum number of electrons in a given 8 subshell (with / — 0) is two. In a p subshell (l = 1) there can be no more than six electrons. The maximum number of electrons in d and / subshells must... 

Other complications are associated with the partitioning of the core and valence space, which is a fundamental assumption of effective potential approximations. For instance, for the transition elements, in addition to the outermost s and d subshells, the next inner s and p subshells must also be included in the valence space in order to accurately compute certain properties (54). A related problem occurs in the alkali and alkaline earth elements, involving the outer s and next inner s and p subshells. In this case, however, the difficulties are related to core-valence correlation. Muller et al. (55) have developed semiempirical core polarization treatments for dealing with intershell correlation. Similar techniques have been used in pseudopotential calculations (56). These approaches assume that intershell correlation can be represented by a simple polarization of one shell (core) relative to the electrons in another (valence) and, therefore, the correlation energy adjustment will be... 

The alkali and alkaline earth metals have their valence electrons in the s subshells. Groups 13 through 18 have their valence electrons located in the p subshells. The transition elements have their valence electrons in the d subshells, and finally, the lanthanides and actinides have their valence electrons in the f sublevel. 

The solution to this problem and the last rule needed to generate the electron configurations for all the atoms came from a German scientist named Friedrich Hund (1896-1997). Hund s rule states that an atom with a higher total spin state is more stable than one with a lower spin state. Because electrons with opposite spin states cancel each other, electrons in p orbitals (and other orbitals except for s) will remained unpaired if possible. Thus, two electrons (or three, for that matter) in a p subshell would remained unpaired. So, the sixth electron in carbon-12 must have the same spin as the fifth one. The Pauli exclusion principle then requires that it fill an empty p orbital. 

This chapter will show that only atoms with partially filled shells (i.e. atoms with unpaired electrons) can possess a net magnetic moment in the absence of an external field. Since main group p block) elements have atoms with filled d subshells and tend to form compounds with other p-block elements that result in filled p subshells in accordance with the octet rule, the vast majority of magnetic materials have historically contained transition metal atoms with partially filled d subshells. Nevertheless, some pure organic compounds with free radicals have been found to exhibit ferromagnetic intermolecular interactions, albeit at very low temperamres (several Kelvins). 

Magnetic quantum number (mi) The magnetic quantum number (mi) describes the orbital s orientation in space. For a given I value, ml has integer values from — / to -i-1. In other words, for the p subshell (I = 1), the mi values are -1, 0, and -I-1, hence three orbitals. 

Now that you understand electron configurations, an orbital diagram can be drawn. Orbital diagrams represent the orbital where each electron is located. An arrow is used to represent each electron spinning in a particular direction. Recall that s subshells have one orbital, p subshells have three orbitals, and d subshells... 

It is equally probable that i orbital electrons will be located in any direction about the nucleus. We say that an s orbital is spherically symmetrical. The li orbital is pictured in Figure 4.8(a). Because an electron with = 1 has three possible me values, any p subshell has three orbitals. Each one lies along one of the coordinate axes—y, or z—as shown in Figure 4.8(b). Each p orbital consists of two 3-dimensional lobes centered on one of the axes. An atom has five 3d orbitals, corresponding to the five possible me values (—2, —1,0, -l-1, and -1-2) for a subshell with = 2. Their orientations are shown in Figure 4.8(c). 

How many unpaired electrons are in an atom in the ground state, assuming that all other subshells are either completely full or empty, if its outermost p subshell contains (a) three electrons, (b) five electrons, (c) four electrons ... 

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