PowerPoint to accompany Molecular Shapes Chapter 8 Molecular Geometry and Bonding Theories Figure 8.2 The shape of a molecule plays an important role in its reactivity. By noting the number of bonding and nonbonding electron pairs, we can easily predict the shape of the molecule. What Determines the Shape of a Molecule? Valence Shell Electron Pair Repulsion Theory (VSEPR) Simply put, electron pairs, whether they are bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. The best arrangement of a given number of electron domains is the one that minimises the repulsions among them.
Electron Domains Electron-Domain Geometries This molecule has four electron domains. We can refer to the electron pairs as electron domains. In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom. Therefore, they count as one electron domain. Table 8.1 These are the electron-domain geometries for two through six electron domains around a central atom. Electron-Domain Geometries Molecular Geometries All one must do is count the number of electron domains in the Lewis structure. The geometry will be that which corresponds to that number of electron domains. Figure 8.6 The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is defined by the positions of only the atoms in the molecules, not the nonbonding pairs. Figure 8.6
Molecular Geometries Electron Domain: Linear Within each electron domain, there might be more than one molecular geometry. In this domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is. Table 8.2 Table 8.2 Electron Domain: Trigonal Planar There are two molecular geometries: Trigonal planar, if all the electron domains are bonding Bent, if one of the domains is a nonbonding pair. Electron Domain: Tetrahedral There are three molecular geometries: Tetrahedral, if all are bonding pairs Trigonal pyramidal, if one is a nonbonding pair Bent, if there are two nonbonding pairs Table 8.2 Table 8.2
Electron Domain: Trigonal Bipyramidal Electron Domain: Trigonal Bipyramidal Figure 8.8 There are two distinct positions in this geometry: Axial Equatorial Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry. Electron Domain: Trigonal Bipyramidal There are four distinct molecular geometries in this domain: Trigonal bipyramidal Seesaw T-shaped Linear Electron Domain: Octahedral All positions are equivalent in the octahedral domain. There are three molecular geometries: Octahedral Square pyramidal Square planar Table 8.3 Table 8.3
The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles Nonbonding pairs are physically larger than bonding pairs. Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule. Double and triple bonds place greater electron density on one side of the central atom than do single bonds. Therefore, they also affect bond angles. Figure 8.7 Shapes of Larger Molecules In larger molecules, it makes more sense to talk about the geometry of a particular atom rather than the geometry of the molecule as a whole. Molecular Shape and Molecular Polarity If a molecule possesses polar bonds, it does not mean the molecule as a whole will be polar. Figure 8.11
Molecular Shape and Molecular Polarity Molecules Containing Polar Bonds Figure 8.12 By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule. Figure 8.13 Covalent Bonding and Orbital Overlap We think of covalent bonds forming through the sharing of electrons by adjacent atoms. In such an approach, this can only occur when orbitals on the two atoms overlap. Covalent Bonding and Orbital Overlap Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electronelectron repulsion. Figure 8.14 However, if atoms get too close, the internuclear repulsion greatly raises the energy. Figure 8.15
Hybrid Orbitals sp Hybrid Orbitals To explain geometries, we often assume that the atomic orbitals of an atom mix to form new orbitals called hybrid orbitals. The process of mixing orbitals as atoms approach each other is called hybridisation. Consider beryllium: In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals. But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds. sp Hybrid Orbitals sp Hybrid Orbitals Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals: These two degenerate orbitals would align themselves 180 from each other. These sp hybrid orbitals have two lobes like a p orbital. One of the lobes is larger and more rounded as is the s orbital. This is consistent with the observed geometry of beryllium compounds: linear. Figure 8.16 Figure 8.17
sp Hybrid Orbitals With hybrid orbitals, the orbital diagram for beryllium would look like this. sp 2 and sp 3 Hybrid Orbitals Using a similar model for boron leads to sp 2 hybridisation, i.e. The sp orbitals are higher in energy than the 1s orbital but lower than the 2p. sp 2 and sp 3 Hybrid Orbitals three degenerate sp 2 orbitals. sp 2 and sp 3 Hybrid Orbitals With carbon we get sp 3 hybridisation, i.e. four degenerate sp 3 orbitals. Figure 8.18
Hybridisation involving d orbitals For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids. Hybridisation involving d orbitals This leads to: - five degenerate sp 3 d orbitals or - six degenerate sp 3 d 2 orbitals. Hybrid Orbitals Once the electrondomain geometry is known, the hybridisation state of the atom is known. Multiple Bonds sigma, - overlap of two orbitals along the internuclear axis. pi, - sideways overlap of two p orbitals. Table 8.4 Figure 8.21
Molecular Orbital (MO) Theory Molecular Orbital of H 2 Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals. Figure 8.22 In H 2 the two electrons go into the bonding molecular orbital. The bond order is half the difference between the number of bonding and antibonding Figure 8.23 electrons, i.e. bond order = ½(no. of bonding electrons - no. of antibonding electrons) Molecular Orbital of H 2 Molecular Orbital of He 2 For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is: In the case of He 2, the bond order would be: 1 2 (2-2) = 0 Figure 8.23 1 2 (2-0) = 1 Therefore, He 2 does not exist. Figure 8.23
Molecular Orbitals of Second-Row Diatomic Molecules Molecular Orbitals for Li 2 and Be 2 The MO diagram looks like this: Figure 8.33 There are both σ and π bonding molecular orbitals and σ* and * antibonding molecular orbitals. Figure 8.27 Molecular Orbitals from 2p Orbitals Molecular Orbitals for B 2 to Ne 2 For atoms with both s and p orbitals, there are two types of interactions: Figure 8.31 Figure 8.26 The s and the p orbitals that face each other overlap in fashion. The other two sets of p orbitals overlap in fashion. The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals. This flips the order of the s and p molecular orbitals in these elements. Figure 8.32
Electron Configurations and Molecular Properties Paramagnetic molecules have one or more unpaired electrons and are attracted into a magnetic field. Diamagnetic molecules have no unpaired electrons and are weakly repelled by a magnetic field. (It is a much weaker effect than paramagnetism.)