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Chapter 10

Chapter 10. Molecular Geometry and Chemical Bonding Theory. Overview. Geometry and Directional Bonding Valence-Shell Electron-Pair Repulsion Theory Dipole Moment and Molecular Geometry Valence Bond Theory Description of Multiple Bonding Molecular Orbital Theory

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Chapter 10

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  1. Chapter 10 Molecular Geometry and Chemical Bonding Theory

  2. Overview • Geometry and Directional Bonding • Valence-Shell Electron-Pair Repulsion Theory • Dipole Moment and Molecular Geometry • Valence Bond Theory • Description of Multiple Bonding • Molecular Orbital Theory • Principles of Molecular Orbital Theory • Electron Configurations of Diatomic Molecules of the Second-Period Elements • Molecular Orbitals and Delocalized Bonding

  3. Molecular Geometry and Directional Bonding • Atoms oriented in very well defined relative positions in the molecule. • Molecular Geometry = general shape of the molecule as determined by the relative positions of the atomic nuclei. • Theories Describing the structure and bonding of molecules are: • VSEPR = considers mostly electrostatics in determining the geometry of the molecule. • Valence Bond Theory = considers quantum mechanics and hybridization of atomic orbitals. • Molecular Orbital Theory = claims that upon bond formation new orbitals that are linear combinations of the atomic orbitals are formed.

  4. Valence Share Electron Pair Repulsion (VSEPR) Theory • Valence Share Electron-Pair Repulsion (VSEPR) model allows us to predict the molecular shape by assuming that the repulsive forces of electron pairs cause them to be as far apart as possible from each other. • Only the valence electron pairs are considered in determining the geometry. • Review: valence electrons are the additional electrons in an atom beyond the inert gas core; for the neutral atom the number is equal to the group number, i.e. C has 4, N has 5, O has 6, and F has 7.

  5. 2 charge clouds, linear 3 charge clouds, trigonal planar 4 charge clouds, tetrahedral 5 charge clouds, trigonal bipyramidal 6 charge clouds, Octahedral Effect of the number of electron pairs around the central atom

  6. PREDICTING EXPECTED GEOMETRY ACCORDING TO VSEPR THEORY. • Lewis dot structure determines the total # of electrons around the central atom. Multiple bonds (double and triple) count as one. • The number of bonding and nonbonding electron pairs determines the geometry of electron pairs and the molecular geometry. • E.g. Determine the geometry of the following: • BeCl2.and CO2 linear • BF3, COCl2, O3, SO2 • CH4, PCl3, H2O • PCl5, SF4, ClF3, XeF2 (lone pair in axial position for a trigonal bipyramidal structure). • SF6, IF5, XeF4 • Lone e Pairs affect geometry more than bonding pairs. • E.g. NH3 has one lone pair of electrons. e pairs repulsion from reduces angle from 109° to 107. H2O with two lone pairs and the angle between the H's and the O is only 105. • Multiple bonds have larger affect on geometry than single bonds: H2C=O (116° instead of 120°); H2C=CH2 (117° instead of 120°).

  7. Figure 9.2 Molecular Shapes 2,3,4 electron pairs.

  8. Figure 9.3 Molecular shapes 5, 6 electron pairs

  9. Polarity of Molecules • Bond dipole a positive charge next to a negative charge. • Dipole moment,  the magnitude of the net bond dipole of a molecule  = Qxr Q = the net charge separation; r = the separation distance. Units: debyes (D) where 1 D = 33.36x1030 Cm. • A polar bond forms when two atoms of between two atoms involved in a bond have significantly different electronegativities. • Most electronegative substance will have a slight negative charge (represented as ) • The positive (electron poor) side of the bond is represented as + or •  points in direction of the negative charge. • Net polarity (dipole moment) of a molecule is obtained using the vector sum of polarities of individual bonds. • E.g. determine if NH3, H2O, CO2 have dipole moments. • E.g.2 determine if either the cis- or trans- isomer of C2H2Cl2

  10. Dipole Moments of Molecules • Dipole moments are easily measured in the laboratory and allow the determination of the net ionization of the molecule. • Complete ionization gives the charge of an electron (1.60x1019 C ) multiplied by r. • the net charge on each of the atoms in a polar bond can be obtained from the measured dipole moment. where n = net charge and r is the radius (in m). E.g. determine the net charge on HCl if the dipole moment is 1.03 D and the bond distance is 127 pm; determine the % ionization. E.g.2 determine the expected dipole moment of NaCl if the bond were completely ionic and then determine the % ionic character. The ionic radii of Na+ and Cl are 102 pm and 181 pm, respectively. Experimental dipole moment is 9.0 D.

  11. MOLECULAR SHAPES:VALENCE BOND THEORY (VBT) • Valence Bond Theory: a quantum mechanical description of bonding that pictures covalent bond formation as the overlap of two singly occupied atomic orbitals. • VSEPR effective but ignores the orbital concepts discussed in quantum mechanics. • H2 forms due to overlap of two 1s orbitals. • Electron densities from p-subshell electrons overlap to produce a bond in F2. • CH4:The 1s orbital of hydrogen must overlap with the 2s and 2p orbitals of carbon. • Presence of electrons from hydrogen adds new waves that are in contact with each other and undergo constructive interference – new waves result. • The s and p orbitals around an atom such as carbon become equivalent and the orbitals become a hybrid (sp3) of the original orbitals. • Hybrid orbitals are as far apart as possible.

  12. Other Kinds of Hybrid Orbitals • Hybridization varies from sp, sp2, up to sp3d2 depending upon the number of orbitals involved in the bonding. • Each of these has a characteristic shape see table in book. • Hybridization determined by using VSEPR to establish the geometry, i.e., the number of electron clouds around the central atom. The number of electron clouds = the number of hybrid orbitals. E.g.: Determine the hybridization of B in BF3. • The  bond formed between an s orbital and a p orbital or even between two p orbitals. • E.g. CH3CH2OH, has all  bonds - even though there are C-C bonds and C-O bonds which each involve the interaction of sp3 orbitals to form the  bonds. • SF6: sp3d2.

  13. VBT: Multiple bonds • C2H4 planar with a trigonal geometry = sp2 hybridization for each of the carbon atoms and they form  bonds with hydrogen. • Each carbon has 4 orbitals in its valence shell. This means one of the p-orbitals for each C is not hybridized. • Proximity to each other results in overlap to give a charge distribution resembling a cloud which is above and below the plane of the molecule and called a  –bond . • Overlap above and below makes rotation of carbon atoms difficult. • E.g. C2H2: sp (linear) hybridized. Leads to the existence of a  bond as well as two  bonds. • Summarizing a • single bond is a  bond, • double bond is a  bond and a  bond, • triple bond is a  bond and 2  bonds. E.g. Dinitrogen difluoride exists as cis and trans isomers( a compound having the same formula with a different arrangement of atoms). Investigate the bonding.

  14. MO Theory of Bonding • Molecular Orbital Theory extends quantum theory and states that electrons spread throughout the molecule in molecular orbitals = region in a molecule in which an electron is likely to be which is similar to the concept discussed in quantum theory. Molecular orbitals are considered to be the result of the combination of atomic orbitals. • Hydrogen: when two atomic orbitals from hydrogen approach each other they form 2 molecular orbitals,  and *, bonding orbital and antibonding orbital respectively. • The energy of the bonding orbital is lower than the original atomic orbital. • The energy of the antibonding orbital is higher than the original atomic orbitals and thus destabilizes the molecule. • The electron distribution of H2 would be: 1s , . An excited state of this molecule would be 1s, .

  15. Fig. 10.34 MO Diagram of N2 Molecular Orbital Theory of Other Diatomic Molecules • He2:  no net stabilization (or bonding). •  a net of one bonding electron. • Bond order: BO = 1/2(nb na) where nb is the number of bonding electrons and na is the number of antibonding electrons. E.g. For He2 BO(He2) = 1/2(2 2) = 0. E.g.2 H2 on the other hand would have a BO(H2) = 1/2(2  0) = 1 or there is a single bond between the two atoms. • Li2:  BO=1/2(4  2) = 1. • Molecule of lithium should be stable. • O2: ; BO= 1/2(10  6) = 2. Last two filled orbitals are antibonding  one elctron in each orbital (Hund’s rule) or two unpaired electrons  O2 a paramagnetic molecule.

  16.                                    MO Levels of 2nd Row Elements Small 2s-2p interaction Large 2s-2p interaction B2 C2 N2 O2 F2 Ne2 Bond Order Magnetic behavior 1 P 2 D 3 D 2 P 1 D 0 P = Paramagnetic; D = Diamagnetic

  17. Delocalized Bonding • Molecular orbital theory handles delocalization quite nicely since molecular orbitals can be said to be spread over the entire. • Metals and energy bands formed by them. • Solidification of metal atoms forms large “molecules” with extensive delocalization of electrons. • Molecular orbitals for all metals are very similar and a continuous band is formed. • They can conduct electricity when the atoms are excited so that an electron occupies an excited state. The energy separation between the occupied and unoccupied orbitals is small so that little energy is required to cause this.

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