Electronic configuration principal quantum number

There is no more room in the 2s orbital for a fifth electron, which appears when we move on to the boron atom. However, another orbital with principal quantum number 2 is available. A 2p orbital accepts the fifth electron, giving the configuration Is ls-lfi. Continuing this process, we obtain the following configurations ... 

To get the electron configuration of ions, a new rule is followed. We first write the electron configuration of the neutral atom. Then, for positive ions, we remove the electrons in the subshell with highest principal quantum number first. Note that these electrons might not have been added last, because of the n + / rule. Nevertheless, the electrons from the shell with highest principal quantum number are removed first. For negative ions, we add electrons to the shell of highest principal quantum number. (That shell has the electrons added last by the n +1 rule.)... 

With Na, the electron configuration of which may also be described as [Ne s1, the third period begins. A similar situation is found for each of the other periods in the Table the number of the period is the principal quantum number of the least tightly bound electron of the first element (an alkali metal) of the period. A few more details of these questions and the characteristics of special points in the Periodic Table are discussed in following paragraphs. The electron configurations of all the elements are given in Chapter 5. 

The block s, on the left of the Table, contains the alkali and alkaline earth metals. Each atom of these metals possesses an inert gas core and one or two electrons in the s orbital of the valence shell, that is, an external electron configuration ns1 or ns2 where n is the value of the principal quantum number, and also the period number in the Periodic Table. Notice however that He, owing to its general chemical inertness and to the behaviour similarity with the other noble gasses is generally placed at the far right of the Table. The p block contains elements corresponding to electron... 

The possible states of electrons are called orbitals. These are indicated by what is known as the principal quantum number and by a letter—s, p, or d. The orbitals are filled one by one as the number of electrons increases. Each orbital can hold a maximum of two electrons, which must have oppositely directed spins. Fig. A shows the distribution of the electrons among the orbitals for each of the elements. For example, the six electrons of carbon (B1) occupy the Is orbital, the 2s orbital, and two 2p orbitals. A filled Is orbital has the same electron configuration as the noble gas helium (He). This region of the electron shell of carbon is therefore abbreviated as He in Fig. A. Below this, the numbers of electrons in each of the other filled orbitals (2s and 2p in the case of carbon) are shown on the right margin. For example, the electron shell of chlorine (B2) consists of that of neon (Ne) and seven additional electrons in 3s and 3p orbitals. In iron (B3), a transition metal of the first series, electrons occupy the 4s orbital even though the 3d orbitals are still partly empty. Many reactions of the transition metals involve empty d orbitals—e.g., redox reactions or the formation of complexes with bases. 

An abbreviated way of presenting electron configuration is to write the principal quantum number and letter of each occupied orbital and then use a superscript to indicate the number of electrons in each orbital. The orbitals of each atom are then written in order of increasing energy levels. For the group 1 elements, this notation is... 

Figure 1.37 illustrates some ionic radii, and Fig. 1.38 shows the relative sizes of some ions and their parent atoms. All cations are smaller than their parent atoms, because the atom loses one or more electrons to form the cation and exposes its core, which is generally much smaller than the parent atom. For example, the atomic radius of Li, with the configuration ls22s, is 157 pm, but the ionic radius of Li+, the bare heliumlike Is2 core of the parent atom, is only 58 pm. This size difference is comparable to that between a cherry and its pit. Like atomic radii, cation radii increase down each group because electrons are occupying shells with higher principal quantum numbers. 

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 ... 

These electrons, which have the biggest principal quantum number, are called valence electrons. Thus we may rewrite the electron configuration of these elements as follows ... 

The order in which the electrons occupy the various n and / states as atomic number increases through the Periodic Table is illustrated in Table 2.1. The prefix number specifies the principal quantum number, the letters s, p, d and / respectively specify the orbitals for which / = 0, 1, 2 and 3, and the superscript specifies the number of electrons in the particular orbital. For brevity the electron configurations for the inert gases are denoted [Ar] for example. 

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