S orbitals filling

Apparent anomalies in the filling of electron orbitals in atoms occur in chromium and copper. In these elements an electron expected to fill an s-orbital fills the d-orbitals instead, (a) Explain why these anomalies occurs, (b) Similar anomalies are known to occur in seven other elements. Using Appendix 2C, identify those elements and indicate for which ones the explanation used to rationalize the chromium and copper electron configurations is valid, (c) Explain why there are no elements in which electrons fill ( / + I )s-orbitals instead of np-orbitals. 

The orbitals being filled for elements in various parts of the periodic table. Note that when we move along a horizontal row (a period), the (n + l)s orbital fills before the nd orbital. The group labels indicate the number of valence electrons (ns plus np electrons) for the elements in each group. 

Size reduction-induced polarization happens in compounds containing N, O, and F and metals with the outermost s-orbit filled with unpaired electrons like Ag, Au, Rh, while happens not to metals with such s-orbit that is empty or filled with paired electrons, such as Pt and Co. Such polarization of substance at the nanoscale creates properties that the bulk counterpart never demonstrates such as the dilute magnetism, catalysis, superhydrophobicity, fluidity, lubricity, as will be addressed in later section. 

The teodeocy to aitaia either a half filled or fully filled set of d orbitals at the expense of the outer s orbital is shown by both chromium and copper and should be noted. This apparent irregularity will be discussed in more detail in Chapter 13. 

We see that the rows of the periodic table arise from filling orbitals of approximately the same energy. When all orbitals of similar energy are full (two electrons per orbital), the next electron must be placed in an s orbital that has a higher principal quantum number, and a new period of the table starts. We can summarize the relation between the number of elements in each row of the periodic table and the available orbitals of approximately equal energy in Table 15-V. 

But I want to return to my claim that quantum mechanics does not really explain the fact that the third row contains 18 elements to take one example. The development of the first of the period from potassium to krypton is not due to the successive filling of 3s, 3p and 3d electrons but due to the filling of 4s, 3d and 4p. It just so happens that both of these sets of orbitals are filled by a total of 18 electrons. This coincidence is what gives the common explanation its apparent credence in this and later periods of the periodic table. As a consequence the explanation for the form of the periodic system in terms of how the quantum numbers are related is semi-empirical, since the order of orbital filling is obtained form experimental data. This is really the essence of Lowdin s quoted remark about the (n + , n) rule. 

These 20 cases do not represent anomalies to the order of orbital filling which is invariably governed by the n + ( rule but are anomalous in the sense that the s orbital is not completely filled before the corresponding d orbital begins to fill. 

There are 2.56 d orbitals available for bond formation. To form 5.78 bonds these would hybridize with the s orbital and 2.22 of the less stable p orbitals. In copper, with one electron more than nickel, there is available an additional 0.39 electron after the hole in the atomic d orbitals is filled. This might take part in bond formation, with use of additional Ap orbital. However, the increase in interatomic distance from nickel to copper suggests that it forms part of an unshared pair with part of the bonding electrons, thus decreasing the effective number of bonds. 

Recently it was pointed out by Zener7 that the atomic moments, in parallel orientation, might react with the electrons in the conduction band in such a way as to uncouple some of the pairs, producing a set of conduction electrons occupying individual orbitals, and with spins parallel to the spins of the atomic electrons. Zener assumed that the conduction band for the transition metals is formed by the 4.s orbitals of the atoms, and that there is somewhat less than one conduction electron per atom in iron, cobalt, and nickel. Like Slater, he attributed the atomic magnetic moments to the partially filled 3d subshell. 

Which a — 2 orbital does the third electron in a lithium atom occupy Screening causes the orbitals with the same principal quantum number to decrease in stability as / increases. Consequently, the 2 S orbital, being more stable than the 2 orbital, fills first. Similarly, 3 S fills before 3 p, which fills before 3 d, and so on. 

The periodic table provides the answer. Each cut in the ribbon of the elements falls at the end of the p block. This indicates that when the n p orbitals are full, the next orbital to accept electrons is the ( + 1 )s orbital. For example, after filling the 3 orbitals from A1 (Z = 13) to Ar (Z = 18), the next element, potassium, has its final electron in the 4 S orbital rather than in one of the 3 d orbitals. According to the aufbau principle, this shows that the potassium atom is more stable with one electron in its 4 orbital than with one electron in one of its 3 (i orbitals. The 3 d orbitals fill after the 4 S orbital is full, starting with scandium (Z = 21). 

A lithium atom has three electrons. The first two electrons fill lithium s lowest possible energy level, the 1. S orbital, and the third electron occupies the 2 5 orbital. The three representations for the ground-state electron configuration... 

Aluminum has Z= 13, so a neutral atom of A1 has 13 electrons. Place the 13 electrons sequentially, using arrows, into the most stable orbitals available. Two electrons fill the = 1 orbital, eight electrons fill the = 2 orbitals, two electrons fill the 3. S orbital, and one electron goes in a 3 orbital. 

For any cation, the empty 4 S orbital is slightly higher in energy than the partially filled 3 d orbital. Thus, the isoelectronic V and Cr cations both have the [Ar] 3 d configuration. On the other hand, the isoelectronic neutral atom scandium has the configuration [Ar]4 3 d ... 

Although electron affinity values show only one clear trend, there is a recognizable pattern in the values that are positive. When the electron that is added must occupy a new orbital, the resulting anion is unstable. Thus, all the elements of Group 2 have positive electron affinities, because their valence S orbitals are filled. Similarly, all the noble gases have positive electron affinities, because their valence a p orbitals are filled. Elements with half-filled orbitals also have lower electron affinities than their neighbors. As examples, N (half-filled 2 p orbital set) has a positive electron affinity, and so does Mn (half-filled 3 d orbital set). 

---