To accommodate these two electron domains, two of the Be atom’s four valence orbitals will mix to yield two hybrid orbitals. There are two regions of valence electron density in the BeCl 2 molecule that correspond to the two covalent Be–Cl bonds. The beryllium atom in a gaseous BeCl 2 molecule is an example of a central atom with no lone pairs of electrons in a linear arrangement of three atoms. In the following sections, we shall discuss the common types of hybrid orbitals. Unhybridized orbitals overlap to form π bonds. Hybrid orbitals overlap to form σ bonds.The type of hybrid orbitals formed in a bonded atom depends on its electron-pair geometry as predicted by the VSEPR theory.All orbitals in a set of hybrid orbitals are equivalent in shape and energy.The number of hybrid orbitals in a set is equal to the number of atomic orbitals that were combined to produce the set. A set of hybrid orbitals is generated by combining atomic orbitals.Hybrid orbitals have shapes and orientations that are very different from those of the atomic orbitals in isolated atoms.They are formed only in covalently bonded atoms. Hybrid orbitals do not exist in isolated atoms.The following ideas are important in understanding hybridization: This description is more consistent with the experimental structure.
(b) Two of the hybrid orbitals on oxygen contain lone pairs, and the other two overlap with the 1 s orbitals of hydrogen atoms to form the O–H bonds in H 2O. The observed angle of 104.5° is experimental evidence for which quantum-mechanical calculations give a useful explanation: Valence bond theory must include a hybridization component to give accurate predictions.įigure 5.7 (a) A water molecule has four regions of electron density, so VSEPR theory predicts a tetrahedral arrangement of hybrid orbitals. Consequently, the overlap of the O and H orbitals should result in a tetrahedral bond angle (109.5°). The valence orbitals in an oxygen atom in a water molecule differ they consist of four equivalent hybrid orbitals that point approximately toward the corners of a tetrahedron ( Figure 5.7). The valence orbitals in an isolated oxygen atom are a 2 s orbital and three 2 p orbitals. The new orbitals that result are called hybrid orbitals. This process of combining the wave functions for atomic orbitals is called hybridization and is mathematically accomplished by the linear combination of atomic orbitals, LCAO, (a technique that we will encounter again later). When atoms are bound together in a molecule, the wave functions combine to produce new mathematical descriptions that have different shapes. The mathematical expression known as the wave function, ψ, contains information about each orbital and the wavelike properties of electrons in an isolated atom. Quantum-mechanical calculations suggest why the observed bond angles in H 2O differ from those predicted by the overlap of the 1 s orbital of the hydrogen atoms with the 2 p orbitals of the oxygen atom. This is not consistent with experimental evidence. The prediction of the valence bond theory model does not match the real-world observations of a water molecule a different model is needed.įigure 5.6 The hypothetical overlap of two of the 2 p orbitals on an oxygen atom (red) with the 1 s orbitals of two hydrogen atoms (blue) would produce a bond angle of 90°. Experimental evidence shows that the bond angle is 104.5°, not 90°. If this were the case, the bond angle would be 90°, as shown in Figure 5.6, because p orbitals are perpendicular to each other. Valence bond theory would predict that the two O–H bonds form from the overlap of these two 2 p orbitals with the 1 s orbitals of the hydrogen atoms. Oxygen has the electron configuration 1 s 22 s 22 p 4, with two unpaired electrons (one in each of the two 2 p orbitals). As an example, let us consider the water molecule, in which we have one oxygen atom bonding to two hydrogen atoms. However, to understand how molecules with more than two atoms form stable bonds, we require a more detailed model. Thinking in terms of overlapping atomic orbitals is one way for us to explain how chemical bonds form in diatomic molecules.
By the end of this section, you will be able to:
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