F Subshell

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. 

Fig. 6 presents the data for citric acid complexes and an extraction of ions, plotted as recommended by Fidelis etSiekerski (23),Peppard (24) and Sinha (25), who introduced different names for the deviations from a smooth log/ = f(Z) variation which occur for a quarter, a half and three-quarters of a filled f subshell. Fig. 6 effectively illustrates this accidents for the 4 f 3-4, 4 f7,4 fIdeviations from a smooth variation) is believed to be a thermodynamic consequence of the nephelauxetic effect. 

The maximum number of electrons in the fourth and subsequent shells is 32. How many electrons will be left after the s, p and d orbitals are filled These electrons will occupy the f subshell. 

You should be able to calculate the number of f orbitals in the f subshell. Remember that each orbital can hold no more than two electrons. The f orbitals are even more complicated than the d orbitals, but you do not have to know or recognise their shapes. 

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. 

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. 

Subshell One or more orbitals in an electron shell of an atom. Subshells are designated s, p, d, and f. The s, p, d, and f subshells contain a maximum of 2, 6, 10, and 14 electrons, respectively (compare with orbital and principal quantum number). 

All orbitals with a value of 1 = 3 are orbitals of the f subshell. The shape of the f orbitals is much more complex than the shape of the s, p and d orbitals. Each subshell has seven orbitals, each with specific orientations. 

Even if the shapes of the orbitals for the first four subshells are given, the shapes of the d and f subshell orbitals are so detailed at this point that they are given to students who are especially interested in learning more about orbital geometry. Look at the figure to the left. 

In contrast, the valence d and f orbitals in heavy atoms are expanded and destabilized by the relativistic effects. This is because the contraction of the s orbitals increases the shielding effect, which gives rise to a smaller effective nuclear charge for the d and f electrons. This is known as the indirect relativistic orbital expansion and destabilization. In addition, if a filled d or f subshell lies just inside a valence orbital, that orbital will experience a larger effective nuclear charge which will lead to orbital contraction and stabilization. This is because the d and f orbitals have been expanded and their shielding effect accordingly lowered. 

The calculations indicate that the Ss subshell should fill at elements 119 and 120, thus making these an alkali and alkaline earth metal, respectively. Next, the calculations point to the filling, after the addition of a Id electron at element 121 of the inner 5g and 6 f subshells, 32 places in all, which the author has termed the superactinide elements and which terminates at element 153. This is followed by the filling of the Id subshell (elements 154 through 162) and 8p sub shell (elements 163 through 168). 

The errors seen in the above examples are of course the result of the necessary incompleteness of orbital and configiaration basis sets. The power of these expansion approaches is that if one works hard enough (uses a sufficiently complete/ or at least appropriate/ basis) one should get the right answer. The recent extensive transition metal hydride studies indicate the possibilities (25-30). Nevertheless/ heavy atom electron correlation involving d and even f subshells is such an enormous paroblem that every alternative should be explored. 

Between these two blocks of elements there are two further blocks containing the transition elements. Strictly speaking, the term transition element applies to an element with a partly filled d or f subshell and so excludes those with d or d and F or electron configurations. However, it is convenient to include copper, silver and gold in this classification as these elements commonly form ions with partly filled d subshells. Although their neutral atoms have d electron configurations, it is the chemistry of their ions which is of primary interest here. Similar arguments apply to ytterbium and nobelium. Their atoms have P s ... 

Although ionization plays a dominant role in the chemistry of the transition elements, the reverse process of adding an electron to their atoms also contributes to their chemical properties. In fact, adding an electron to the valence shell of most transition elements is an exothermic process. This might be anticipated for elements in which partly filled d or f subshells are present. However, for zinc, cadmium and mercury, which have filled valence shells [ d ( +l)s (n = 3, 4 or 5)], the process of electron addition is endothermic. 

Ge, As, Se, Br, and Kr). Zinc, cadmium, and mercury are often classified as main group elements. The periodic table is divided into blocks. The s-block elements have valence configuration si or s2. The p-block elements have valence configuration slpl to s2p6. The d-block and /-block elements usually have two electrons in the outermost s-orbital but have partially filled d or f subshells in an inner orbital. 

The n = 4 principal shell contains four subshells, an s-, p-, d-, and an f-subshell. As before, they are known as the 4s-subshell, 4p-subshell, 4d-subshell, and the 4f-subshell. 

The orbital is a region of space where an electron assigned to that orbital is most likely to be found. Each orbital can hold a maximum of two electrons. The subshells are composed of one or more orbitals. There is one orbital in an s-subshell. There are three orbitals in a p-subshell, five in a d-subshell, and seven in an f-subshell. The number of orbitals in a principal shell equals n2. There are nine orbitals in the n = 3 principal shell, 32 = 9. 

An f-subshell can hold a maximum of 14 electrons, 2 in each of the 7 orbitals. 

In p-, d-, and f-subshells that have three, five, or seven orbitals, each orbital is filled singly before any orbital contains two electrons. This is Hund s rule. 

The orbitals in p-, d-, or f-subshells must be completely filled with electrons before moving to the next higher subshell. A p-subshell can hold a maximum of six electrons a d-subshell, ten electrons an f-subshell, fourteen. 

An s subshell is spherically symmetric and can hold a maximum of 2 electrons. A p subshell is dumbbell-shaped and holds 6 electrons, a d subshell 10 electrons, and an f subshell 14 electrons, with increasingly complicated shapes. 

Diatomic hydrides are a special case, since H has only a Is valence AO. Consider HF as an example. The ground-state configurations of the atoms are Is for H and ls 2s 2p for F. We expect the filled Is and 2s F subshells to take little part in the bond-... 

Figure 6.31 allow us to reexamine the concept of valence electrons. Notice, for example, that as we proceed from Cl ([Ne]3s 3p ) to Br ([Arj3d 4s 4p ) we add a complete subshell of 3d electrons to the electrons beyond the [Ar] core. Although the 3d electrons are outer-shell electrons, they are not involved in chemical bonding and are therefore not considered valence electrons. Thus, we consider only the 4s and 4p electrons of Br to be valence electrons. Similarly, if we compare the electron configurations of Ag (element 47) and Au (element 79), we see that Au has a completely full 4/ subsheU beyond its noble-gas core, but those 4/electrons are not involved in bonding. In general,/or representative elements we do not consider the electrons in completely filled d or f subshells to be valence electrons, and for transition elements we do not consider the electrons in a completely filled f subshell to be valence electrons. 

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