Hybrid orbitals VSEPR theory

Now that we know how to determine hybridization states, we need to know the geometry of each of the three hybridization states. One simple theory explains it all. This theory is called the valence shell electron pair repulsion theory (VSEPR). Stated simply, all orbitals containing electrons in the outermost shell (the valence shell) want to get as far apart from each other as possible. This one simple idea is all you need to predict the geometry around an atom. First, let s apply the theory to the three types of hybridized orbitals. 

The Cl—F and Cl—Cl bonds in the cation are then formed by the overlap of the half-filled sp3 hybrid orbitals of the central chlorine atom with the half-filled p-orbitals of the terminal Cl and F atoms. Thus, by using sp3 hybridization, we end up with the same bent molecular geometry for the ion as that predicted by VSEPR theory (when the lone pairs on the central atom are ignored)... 

The VSEPR theory is only one way in which the molecular geometry of molecules may be determined. Another way involves the valence bond theory. The valence bond theory describes covalent bonding as the mixing of atomic orbitals to form a new kind of orbital, a hybrid orbital. Hybrid orbitals are atomic orbitals formed as a result of mixing the atomic orbitals of the atoms involved in the covalent bond. The number of hybrid orbitals formed is the same as the number of atomic orbitals mixed, and the type of hybrid orbital formed depends on the types of atomic orbital mixed. Figure 11.7 shows the hybrid orbitals resulting from the mixing of s, p, and d orbitals. 

In order for VSEPR theory to make sense, it must be combined with another idea hybridization. Hybridization refers to the mixing of atomic orbitals into new, hybrid orbitals of equal energy. Electron pairs occupy equivalent hybrid orbitals. It s important to realize that the hybrid orbitals cire all equivalent, because that helps you understand the shapes that emerge from the electron pairs trying to distance themselves from one another. If electrons in a pure p orbital are trying to distance themselves from electrons in another p orbital and from electrons in an s orbital, the resulting shape may not be symmetrical, because s orbitals cire different from p orbitals. But if all these electrons occupy identical hybrid orbitals (each orbital is a little bit s and a little bit p), then the resulting shape is more likely to be symmetrical. 

Real molecules have all sorts of symmetrical shapes that just don t make sense if electrons truly occupy only pure orbitals (like s and p). The mixing of pure orbitals into hybrids allows chemists to explain the symmetrical shapes of real molecules with VSEPR theory. This kind of mixing must in some sense actually occur, as the case of methane, CH, makes clear. 

The single 2s orbital combines with the three 2p orbitals to create four identical sp hybrid orbitals. The fact that each sp orbital is identical is important because VSEPR theory can now explain the symmetrical shape of methane the tetrahedron. 

Although this phenomenon represents an exception to the rules, it s somewhat less annoying than other exceptions because hybridization allows for the nicely symmetrical orbital geometries of actual atoms within actual molecules. VSEPR theory presently clears its throat to point out that the negative charge of the electrons within the hybridized orbitals causes those equivalent orbitals to spread as far apart as possible from one another. As a result, the geometry of sp -hybridized methane (CH ), for example, is beautifully tetrahedral. 

Table 5.1 also includes the types of hybridization for each of the basic shapes which are used in VSEPR and valence bond theories. Hybridization, or mixing of the indicated atomic orbitals, in each case produces hybrid orbitals that radiate from the central atom to where the ligand atoms are situated. This is done to ensure that there are localized... 

The VSEPR theory assumes that the four electrons from the valence shell of the carbon atom plus the valency electrons from the four hydrogen atoms form four identical electron pairs which, at minimum repulsion, give the observed tetrahedral shape. To rationalize the tetrahedral disposition of four bond-pair orbitals with those of the 2s and three 2p atomic orbitals of the carbon atom, sp3 hybridization is invoked. 

This chapter reviews molecular geometry and the two main theories of bonding. The model used to determine molecular geometry is the VSEPR (Valence Shell Electron Pair Repulsion) model. There are two theories of bonding the valence bond theory, which is based on VSEPR theory, and molecular orbital theory. A much greater amount of the chapter is based on valence bond theory, which uses hybridized orbitals, since this is the primary model addressed on the AP test. 

The hybridized orbital approach is a simplified way of predicting the geometry of a molecule with three or more atoms by mixing the valence orbitals of its central atom. An alternative approach, valence shell electron-pair repulsion (VSEPR) theory, accomplishes the same thing in a more qualitative way. 

There is a close relationship between the VSEPR theory discussed in Section 3.9 and the hybrid orbital approach, with steric numbers of 2, 3, and 4 corresponding to sp, sp, and sp hybridization, respectively. The method can be extended to more complex structures (fsp hybridization (see Sec. 8.7), which gives six equivalent hybrid orbitals pointing toward the vertices of a regular octahedron, is applicable to molecules with steric number 6. Both theories are based on minimizing the energy by reducing electron-electron repulsion. 

Both nitrogen atoms have steric number 4 and are sp hybridized, with H—N—H and H—N—N angles of approximately 109.5°. The extent of rotation about the N—N bond cannot be predicted from the VSEPR theory or the hybrid orbital model. The full three-dimensional structure of hydrazine is shown in Figure 6.46. 

Valence Shell Electron Pair Repulsion (VSEPR) Theory Hybridization of Atomic Orbitals, sp, sp, sp Single Bonds Conformational Isomers Pi Bonds Pi Barrier to Rotation C/s and Trans, 2p-3p Triple Bonds Cumulenes... 

To be consistent with experimental findings and the predictions of the VSEPR theory, the VB theory must explain three equivalent B—F bonds. Again we use the idea of hybridization. Now the 2s orbital and two of the 2p orbitals of B hybridize to form a set of three equivalent sp hybrid orbitals. 

Each C atom in C2Hg has four regions of high electron density. The VSEPR theory tells us that each C atom has tetrahedral electronic geometry the resulting atomic arrangement around each C atom has one C and three H atoms at the corners of this tetrahedral arrangement. The VB interpretation is that each C atom is sp hybridized. The C—C bond is formed by overlap of a half-filled sp hybrid orbital of one C atom with a half-filled sp hybrid orbital of the other C atom. Each C—H bond is formed by the overlap of a half-filled sp hybrid orbital on C with the half-filled Ir orbital of an H atom. 

We must remember that theory (and its application) depends on fact, not the other way around. Sometimes the experimental facts are not consistent with the existence of hybrid orbitals. In PH3 and ASH3, each H—P—H bond angle is 93.7°, and each H—As—H bond angle is 91.8°. These angles very nearly correspond to three/ orbitals at 90° to each other. Thus, there appears to be no need to use the VSEPR theory or hybridization to describe the bonding in these molecules. In such cases, we just use the pure atomic orbitals rather than hybrid orbitals to describe the bonding. 

In all five of these hybrid orbital schemes, the use of hybridisation is only to give an improved directional overlap of orbitals to form two electron pair covalent bonds. Hybridisation does not determine the basic stereochemistry. This must still be determined by VSEPR theory and only then can hybridisation schemes be invoked to describe, more effectively, the covalent bonding present. These hybridisation schemes may equally be applied to cations and anions. The NH4 cation and BF4" anion have already been shown to involve a tetrahedral stereochemistry (Figure 6.4, examples 3 and 4) consequently the bonding in both ions may be described as involving sp hybridisation. 

Given the basic shapes of mononuclear molecules, cations and anions as determined by VSEPR theory, the bonding involved can then be described by the various types of hybrid orbitals, including double and triple bonding. In polynuclear molecules, VSEPR theory can be used to determine the stereochemistry at the separate atom centres present. Consequently the bonding at these separate atom centres can still be described by the appropriate types of hybrid orbitals, including multiple bonding. 

Theory is a term that is very widely used by chemists. To take the area of chemical bonding as an example, chemists widely refer to molecular orbital (hereafter MO) theory, valence bond (VB) theory, hybridization theory, valence shell electron pair repulsion (VSEPR) theory, and ligand field theory. And even those probably do not exhaust the list. 

In HF the F 2s orbital is too low in energy to be involved significantly in bonding, but in BH both 2s and 2p orbitals on B can contribute to MOs with H Is. This situation is described as sp hybridization, and leads to bonding and nonbonding MOs with a spatial localization similar to that assumed in VSEPR theory. 

Valence bond theory is one of the two quantum mechanical approaches that explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of overlapping atomic orbitals. Using the concept of hybridization, valence bond theory can explain molecular geometries predicted by the VSEPR model. However, the assumption that electrons in a molecule occupy atomic orbitals of the individual atoms can only be an approximation, since each bonding electron in a molecule must be in an orbital that is characteristic of the molecule as a whole. 

These questions and other aspects concerning energy considerations, molecular orbital theory, Lewis structures, VSEPR theory and orbital hybridization theory will be answered and covered in this section. 

By use of VSEPR theory, in which the Lewis structures helped us determining the number of surrounding electron groups, we are now able to predict actual structures of molecules and composite ions. However from the VSEPR theory we know nothing about the chemical bond itself Where are the bond electrons actually placed Or more specific In which types of orbitals are the bond electrons placed The answer to this can be found in the orbital hybridization theory which is the topic in the next section. 

However the orientations of the atomic orbitals in space do not fit the directions predicted by the VSEPR theory according to Table 2- 1 on page 70. For this reason other orbitals than the atomic orbitals must be present in the molecules and composite ions in order to give the right bond directions according to the VSEPR theory. These orbitals are a type of molecular orbitals (also mentioned in section 2.2.2 Molecular orbital theory) which are called hybrid orbitals. These hybrid orbitals thus host the valence electrons which constitutes the chemical bond between the atoms. 

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