![]() ![]() The dots indicate the locations of the nuclei. These bonds are described in more detail later in this chapter.\): Sigma (σ) bonds form from the overlap of the following: (a) two s orbitals, (b) an s orbital and a p orbital, and (c) two p orbitals. In any multiple bond, there will be one σ bond and the remaining one or two bonds will be π bonds. Between any two atoms, the first bond formed will always be a σ bond, but there can only be one σ bond in any one location. The double bond consists of one σ bond and one π bond, and the triple bond consists of one σ bond and two π bonds. As the Lewis structures below suggest, O 2 contains a double bond, and N 2 contains a triple bond. While all single bonds are σ bonds, multiple bonds consist of both σ and π bonds. The dots indicate the location of the nuclei. ![]() Pi (π) bonds form from the side-by-side overlap of two p orbitals. Along the axis itself, there is a node, that is, a plane with no probability of finding an electron. In a π bond, the regions of orbital overlap lie on opposite sides of the internuclear axis. The dots indicate the locations of the nuclei.Ī pi bond (π bond) is a type of covalent bond that results from the side-by-side overlap of two p orbitals, as illustrated in Figure. Sigma (σ) bonds form from the overlap of the following: (a) two s orbitals, (b) an s orbital and a p orbital, and (c) two p orbitals. Single bonds in Lewis structures are described as σ bonds in valence bond theory. A σ bond is a covalent bond in which the electron density is concentrated in the region along the internuclear axis that is, a line between the nuclei would pass through the center of the overlap region. The overlap of two s orbitals (as in H 2), the overlap of an s orbital and a p orbital (as in HCl), and the end-to-end overlap of two porbitals (as in Cl 2) all produce sigma bonds (σ bonds), as illustrated in Figure. Hence, a bond can form by the pairing of each hydrogen electron with an oxygen electron and the overlap of the orbitals they occupy. (b) Any other arrangement results in less overlap. (a) The overlap of two p orbitals is greatest when the orbitals are directed end to end. Figure 1 illustrates this for two p orbitals from different atoms the overlap is greater when the orbitals overlap end to end rather than at an angle. Greater overlap is possible when orbitals are oriented such that they overlap on a direct line between the two nuclei. In addition to the distance between two orbitals, the orientation of orbitals also affects their overlap (other than for two s orbitals, which are spherically symmetric). Orbitals that overlap extensively form bonds that are stronger than those that have less overlap. The strength of a covalent bond depends on the extent of overlap of the orbitals involved. The mutual attraction between this negatively charged electron pair and the two atoms’ positively charged nuclei serves to physically link the two atoms through a force we define as a covalent bond. According to valence bond theory, a covalent bond results when two conditions are met: (1) an orbital on one atom overlaps an orbital on a second atom and (2) the single electrons in each orbital combine to form an electron pair. We say that orbitals on two different atoms overlap when a portion of one orbital and a portion of a second orbital occupy the same region of space. Valence bond theory describes a covalent bond as the overlap of half-filled atomic orbitals (each containing a single electron) that yield a pair of electrons shared between the two bonded atoms. In the following sections, we will discuss how such bonds are described by valence bond theory and hybridization. One popular theory holds that a covalent bond forms when a pair of electrons is shared by two atoms and is simultaneously attracted by the nuclei of both atoms. A more complete understanding of electron distributions requires a model that can account for the electronic structure of molecules. When atoms bond to form molecules, atomic orbitals are not sufficient to describe the regions where electrons will be located in the molecule. However, these predictions only describe the orbitals around free atoms. We can use quantum mechanics to predict the specific regions around an atom where electrons are likely to be located: A spherical shape for an s orbital, a dumbbell shape for a p orbital, and so forth. There are successful theories that describe the electronic structure of atoms. ![]()
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