Filling the 2p Orbitals

The electron configuration of carbon, atomic number 6, is ls 2s 2p. The four outer (valence) electrons are shown by the Lewis symbol 

Two of the four valence electrons are represented by a pair of dots,to indicate that these electrons are paired in the same orbital. These are the two 2s electrons. The other two are shown as individual dots to represent two unpaired 2p electron in separate orbitals. 

The next element to be considered is oxygen, which has an atomic number of 8. Its electron configuration is ls 2s 2p. It is the first element in which it is necessary for 2 electrons to occupy the same p orbital, as shown by the following orbital diagrams  

The electron configuration of fiuorine, atomic number 9, is s 2s 2p. Its Lewis symbol 

atomic number 10, is at the end of the second period of the periodic table and is a noble gas, as shown by its Lewis symbol 

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The choice of filling the 2p orbitals alphabetically is arbitrary as they are three-fold degenerate. 

Period 2 begins with lithium, which has three electrons. The first two electrons fill the I5 orbital, while the third electron goes into the 2s orbital, the sublevel with the next lowest energy. In beryllium, another electron is added to complete the 2s orbital. The next six electrons are used to fill the 2p orbitals. The electrons are added one at a time from boron to nitrogen, which gives three half-filled 2p orbitals. Because orbitals in the same sublevel are equal in energy, there is less repulsion when electrons are placed in separate orbitals. 

Section 2 21 Carbon is sp hybridized m acetylene and the triple bond is of the ct + Tt + Tt type The 2s orbital and one of the 2p orbitals combine to give two equivalent sp orbitals that have their axes m a straight line A ct bond between the two carbons is supplemented by two tr bonds formed by overlap of the remaining half filled p orbitals... 

Next, we half-fill the lone unhybridized 3p orbital on sulfur and the lone 2p orbital on the oxygen atom with a formal charge of zero (atom B). Following this, the 2p orbital of the other two oxygen atoms (atoms C and D), are filled and then lone pairs are placed in the sp2 hybrid orbitals that are still empty. At this stage, then, all 24 valence electrons have been put into atomic and hybrid orbitals on the four atoms. Now we overlap the six half-filled sp2 hybrid orbitals to generate the cr-bond framework and combine the three 2p orbitals (2 filled, one half-filled) and the 3p orbital (half-filled) to form the four 7t-molecular orbitals, as shown below ... 

To form a bond, the filled 2s orbital and one of the 2p orbitals combine and give two half-filled sp orbitals. 

Hund s rule explains that when degenerate orbitals (orbitals that have same energy) are present but not enough electrons are available to fill all the shell completely, then a single electron will occupy an empty orbital first before it will pair up with another electron. This is understandable, as it takes energy to pair up electrons. Therefore, the six electrons in the carbon atom are filled as follows the first four electrons will go to the Is and 2s orbitals, a fifth electron goes to the 2px, the sixth electron to the 2py orbital and the 2p orbital will remain empty. 

The layout of the periodic table (Fig. 2.5) reflects the shell structure of the electrons. Hydrogen and helium have only -shell electrons. The elements in row two have and L-shell electrons, with the Is orbitals always filled and the 2s and 2p orbitals filled in succession. Those in row three have and L-shell electrons, with Is, 2s, and 2p orbitals filled, and the 3 s and 3p orbitals are filled in succession. Elements in the fourth row have K, L, and M-shell electrons, with the Is, 2s, 3s, 2p, and 3p orbitals completely filled. After the 4s orbitals are filled, the 3d orbitals are filled, giving the transition metals. Then come the 4p orbitals. Row five is filled in an analogous fashion. In row six, the lanthanides, which fit between lanthanum and hafnium, reflect the appearance of the N-shell electrons, which fill the f orbitals. Row seven, which contains the actinides, also has K, L, M, and N-shell electrons. 

A boron atom, B (atomic number 5), in its lowest energy state has four of its five electrons filling the Is and 2s orbitals. Its fifth electron may reside in any one of the 2p orbitals, all of which are at the same energy level ... 

Fig. 7.21 Bonding of a mole of lithtum atom 2s orbitals to form a half-filled band. Heavy shading indicates the filled portion of the band, the top of which is called the Fermi level.

The classification of bonds as single, double or triple is not always clear-cut. For example, the Si-F bond may have some double-bond character, depending on the extent of overlap between empty silicon 3d orbitals and filled F 2p orbitals, and this will vary from one situation to another. The large range quoted for the C(sp2)-C(sp2) bond energy reflects the variable amount of p -pn overlap which may be present. 

Electrons are added to orbitals in order of increasing energy, filling the l.v orbital before any electrons occupy the 2s level. The 2s orbital is filled before any of the 2p orbitals, and the 3s orbital is filled before any of the 3p orbitals. All the 2p orbitals (2px, 2py, 2pz) are of equal energy, and each is singly occupied before any is doubly occupied. The same holds for the 3p orbitals. ... 

Figure 8.21 shows schematically a set of Is-, 2s, 2p and 35 core orbitals of an atom lower down the periodic table. The absorption of an X-ray photon produces a vacancy in, say, the Is orbital to give A+ and a resulting photoelectron which is of no further interest. The figure then shows that subsequent relaxation of A+ may be by either of two processes. X-ray fluorescence (XRF) involves an electron dropping down from, say, the 2p orbital to fill the Is... 

These numbers explain the shape of the periodic table. Each element has one more electron (and one more proton and perhaps more neutrons) than the one before. At first the lowest energy shell (n = 1) is filled. There is only one orbital, Is, and we can put one or two electrons in it. There are therefore two elements in this block, H and He. Next we must move to the second shell ( = 2), filling 2s first so we start the top of groups 1 and 2 with Li and Be. These occupy the top of the red stack marked s block because all the elements in this block have one or two electrons in their outermost s orbital and no electrons in the outermost p orbital. Then we can start on the 2p orbitals. There are three of these so we can put in six electrons and get six elements B, C, N, O, F, and Ne. They occupy the top row of the black p block. Most of the elements we need in this book are in those blocks. Some, Na, K, and Mg for example, are in the s block and others, Si, P, and S for example, are in the second row of the p block. 

They will fill orbitals with the lowest energy first. The Is orbital will fill before the 2s orbital, which will fill before the 2p orbitals. 

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