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how atoms can combine their atomic orbitals to form molecular orbitals. We need to extend our discussion of molecular orbitals to learn how more complicated molecules are built up from atoms. First, we will consider other diatomic molecules, for example, N2, O2, F2, and HF. N2, O2, and F2 (nitrogen, oxygen, and fluorine) are called homonuclear diatomics because the two atoms are the same. HF (hydrogen fluoride) is a heteronuclear diatomic because the two atoms are different. Examination of homonuclear diatomics will let us go beyond the hydrogen molecule, which is a special case. The nature of molecular orbitals for heteronuclear diatomics is an important step toward the study of polyatomic molecules that compose most of the molecular substances that surround us, from alcohol to fats.

The hydrogen molecule is the only neutral molecule that only uses electrons in its 1s orbitals to form a chemical bond. The electrons that an atom employs in bonding are called the valence electrons. In the molecules N2, O2, F2, and HF, 2s and 2p orbitals are involved in bonding. The 2s and 2p electrons are the valence electrons. N, O, and F are in the second row of the Periodic Table. Atoms in the third row of the Periodic Table, such as P, S, and Cl (phosphorus, sulfur, and chlorine), will have 3s and 3p valence electrons involved in bonding. Atoms in the third and higher rows of the Periodic Table can also employ d electrons in forming chemical bonds. Here we will focus on the second row elements. The second row elements are of great importance, and the ideas that will be developed are sufficiently general to cover the nature of chemical bonding of heavier elements.

SIGMA (σ) AND PI (π) BONDS

As shown in Figure 12.2, when two hydrogen atoms form the H2 molecule, the two hydrogen 1s orbitals combine to form the bonding molecular orbital. There is electron density along the line that connects the nuclei. A σ (sigma) bonding or antibonding molecular orbital has electron density along the line connecting the nuclei. In H2, we say that a σ bond is formed usingaσbonding MO. s orbitals always form σ bonds. There is no way to bring together two s orbitals and not have electron density along the line connecting the nuclei. This is not true for p orbitals.

Because of their shapes, there are two ways for a pair of p orbitals to come together, as is illustrated in Figure 13.1. The drawings in Figure 13.1 show the p orbitals very schematically. These are actually probability amplitude waves that have a diffuse probability distribution of finding the electron some distance from the nucleus. Here the outline gives a representation of the general shape of a p orbital. Better illustrations are shown in Figure 10.7. Recall that a p orbital has a nodal plane between the two lobes. A nodal plane is a plane in which the probability of finding the electron is zero. For a pz orbital, the nodal plane is the xy plane (see Figure 10.7). The probability of finding an electron in some region of space is frequently called the electron density. A high density means there is a high probability of finding the electron.

FIGURE 13.1. A pair of p orbitals brought close to each other. Upper portion: the orbitals are brought together end to end. There is electron density along the line connecting the nuclei. Bottom portion: the orbitals are brought together side to side. There is no electron density along the line connecting the nuclei.

The upper portion of Figure 13.1 shows two p orbitals brought close to each other end to end. The lobes are pointing at each other. The nuclei are represented by the dots. The dashed line is the line connecting the nuclei. There is clearly electron density along the line connecting the two nuclei. The lower portion of the figure shows the two p orbitals brought close to each other side to side. The nodal plane is perpendicular to the plane of the page. The nuclei lie in the nodal plane. There is no electron density along the line connecting the nuclei. The lobes of the p orbitals have signs. One lobe is positive and the other