D subshell

A transition element is an element whose atom has an incomplete d subshell, or which gives rise to a cation or cations with an incomplete d subshell. 

The incompletely filled d-subshell is responsible for the wide range of colors shown by compounds of the d-block elements. Furthermore, many d-metal compounds are paramagnetic (see Box 3.2). One of the challenges that we face in this chapter is to build a model of bonding that accounts for color and magnetism in a unified way. First, though, we consider the physical and chemical properties of the elements themselves. 

In bivalent nickel, palladium, or platinum there are eight electrons in the outer d subshell. Putting them two to an orbital, they occupy a minimum of four of the five d orbitals, leaving only one d orbital available for bond formation through combination with s, px, py, and pz of the next outer shell. It is found that only four strong bond orbitals can be formed. These four lie in a plane and are directed towards the four comers of a... 

The elements may be divided into types (Fig. 17-10), according to the position of the last electron added to those present in the preceding element. In the first type, the last electron added enters the valence shell. These elements are called the main group elements. In the second type, the last electron enters a d subshell in the next to last shell. These elements are the transition elements. The third type... 

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. The maximum number of electrons in any outermost shell (except the first shell) is 8. The fifth shell starts before the d subshell of the fourth shell starts. 

Ans. (a) 10. This is a d subshell, with five orbitals corresponding to rn values of - 2, - 1, 0, 1, 2. Each orbital can hold a maximum of 2 electrons, and so the subshcll can hold 5x2= 10 electrons, ib) 2. This is an s subshell, (c) 6. This is a p subshcll (with three orbitals), (d) 2. This is an s subshcll. The principal quantum number docs not matter. [Compare part (/>).]... 

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. 

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. 

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. 

The d block transition metals are metals with an incomplete d subshell in at least one of their ions. We consider the first row of the transition metals as being from scandium to zinc and the second row from yttrium to cadmium. Platinum and gold are in the third row. Most of the common metals in everyday use are transition metals. 

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. 

Scandium always forms Sc + ions and, because the 4s electrons are lost first, the electronic configuration of the Sc + ion is Is 2s 2p 3s 3p so there are no electrons in the d subshell. 

There is a special stability associated with all the orbitals in the d subshell being half-filled and so Fe " is more stable than Fe. ... 

The third shell or energy level is M and may contain a maximum of 18 electrons its orbital is called the d subshell, and it may have a maximum of 10 electrons for example,... 

There is an important class of rearrangements of strained cyclic a-bonded systems to give less strained ir-bonded qrstems which occur under the influence of transition metal catalysts although the uncatalysed proce is Woodward-Hoffman forbidden and slow. Examples are the conversion of cubanes XXII and bis-homocubanes XXIll to syn-tricyclooctadienes XXIV and related species XXV and of quadricyclene (XXVI) to norbomadiene (XXVII) [Ag, however, converted cubane and related species to the previously unrecognised species cuneane (XXVIII) and its relatives as do some electrophiles with incompletely filled d-subshells ... 

When a d-metal atom loses electrons to form a cation, it first loses its outer s-electrons. However, most transition metals form ions with different oxidation states, because the d-electrons have similar energies and a variable number can also be lost when they form compounds. Iron, for instance, forms Fe2+ and Fe3+ copper forms Cuf and Cu2+. The reason for the difference between copper and potassium, which forms only K+, can be seen by comparing their second ionization energies, which are 1958 kj-mol 1 and 3051 kj-mol-1, respectively. To form Cu2+, an electron is removed from the d subshell of [Ar]3d10 but to form K2+, the electron would have to be removed from potassium s argonlike core. Because such huge amounts of energy are not readily available in chemical reactions, a potassium atom can lose only its 4s-electron. 

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. 

These ions do not have a true noble gas electron configuration because they have an additional filled d subshell. 

The transition elements occupy A the central part of the periodic table, bridging the gap I between the active s-block metals of groups 1A and 2A on the left and the p-block metals, semimetals, and nonmetals of groups 3A-8A on the right (Figure 20.1). Because the d subshells are being filled in this region of the periodic table, the transition elements are also called the d-block elements. 

Each d subshell consists of five orbitals and can accommodate 10 electrons, so each transition series consists of 10 elements. The first series extends from scandium through zinc and includes many familiar metals, such as chromium, iron, and copper. The second series runs from yttrium through cadmium, and the third series runs from lanthanum through mercury. In addition, there is a fourth transition series made up of actinium through the recently discovered and as yet unnamed element 112. 

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. 

The remarks in the previous paragraph are particularly relevant to UPS studies of coordination compounds where the central atom has a partly-filled d subshell. This area is restricted by the low volatility and/or thermal stability characteristic of most coordination compounds UPS is... 

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