Chapter 7 Molecular Geometry and Bonding Theories We will not cover Section 7 6 Molecular Orbital Theory General Information There are several bonding theories each builds and expands on the pervious one to explain bonding and to address most of the experimental observations These theories are 1 2 3 4 I A 1 2 3 4 Lewis structures according to the Lewis bonding theory chemical bonds result from the sharing of electron pairs between atoms While this theory explains electron distribution the strength and length of bonds it does not explain the three dimensional geometry of molecules Valence Shell Electron Pair Repulsion VSEPR provides a simple means for predicting the shapes of molecules in terms of electron pair repulsions however it does not explain why bonds form between atoms Valence bond theory explains bonding in terms of overlap of atomic orbitals and the formation of hybrid orbitals It further demonstrates that bonds form because the resulting molecule has a lower potential energy than the isolated atoms Molecular Orbital Theory explains bonding in terms of the formation of delocalized molecular orbitals that extend over several atoms or the whole molecule This theory can predict the magnetic properties of molecules and ions Molecular Geometry General Information Definitions Molecular geometry the three dimensional arrangement of atoms in molecules or ions VSEPR Valence Shell Electron Repulsion Theory explains molecular geometry in terms of repulsive force between terminal atoms The terminal atoms will be as far from each other as possible to minimize repulsive forces Valence shell the outermost shell of an atom that contains the electrons that are involved in bonding Electron domain the number of entities bonds and lone pairs around the central atom Single double and triple bonds are considered to be one domain as is a lone pair 1 B VSEPR Model and its Rules 1 2 3 Double and triple bonds are treated as one domain If a molecule has two or more resonance structures the VSEPR model can be applied to any one of the structures Formal charges are usually not shown Molecules in Which the Central Atom Has No Lone Pairs A central atom B terminal atoms Electron Domain Geometry Geometry Angles Number of E n Pairs Molecular Bond Examples AB2 Linear Linear 180 BeCl 2 AB3 Trigonal planar Trigonal planar 120 AB4 Tetrahedral Tetrahedral 109 5 BF 3 CH 4 PCl AB5 Trigonal bipyramidal Trigonal bipyramidal 120 180 90 5 AB6 Octahedral Octahedral 90 SF 6 2 C Deviation from Ideal Bond Angles Lone pairs on the central atom will affect the bond angles because the lone pairs experience less attraction by the central atom their electron cloud expands and forces the bonding pairs closer to one another 3 2 Examples Use VSEPR to predict the geometry and bond angles in the following species SiBr4 NF3 CS2 NO3 4 IF3 XeF5 D Geometry of Molecules with More Than One Central Atom 1 2 For molecules that have more than one central atom determine the electron domain geometry and molecular geometry around each central atom For example in methyl alcohol the electron domain geometry about the carbon and oxygen atoms are both tetrahedral but the molecular geometry about the oxygen is bent and about the carbon it is tetrahedral H H H C O H than 109 5 the C O H bond angle is 3 Determine the geometry about each central atom in II Molecular Geometry and Polarity General Information 5 1 2 3 4 1 2 3 4 Molecular geometry determines molecular polarity which in turn influences the physical chemical and biological properties of a substance A diatomic molecule is polar if the two atoms that are connected by a chemical bond have different electronegativities H F In more complex molecules polarity is determined not only by the polarity of the bonds but also by the geometry of the molecule The polarity of a molecule can be determined qualitatively by vector addition A completely symmetrical molecule is nonpolar BF3 CH4 SF6 III Valence Bond Theory A Formation of the hydrogen molecule H2 Two individual hydrogen atoms have a specific energy As the two atoms approach and attract one another their potential energy decreases When the distance between the two atoms is about 74 pm their potential energy is at a minimum and the two atoms bond to one another forming the H2 molecule If the two atoms were forced closer than 74pm significant repulsion would occur potential energy would increase and the molecule would break apart There is an would 6 B 1 2 3 C 1 2 3 4 A 1 optimum distance between the two atoms General Information About Valence Bond Theory According to valence bond theory a covalent bond will form between two atoms if the potential energy of the resulting molecule is lower than the combined potential energies of the isolated atoms When a bond forms energy is released Bond formation is exothermic Important features of valence bond theory A bond forms when singly occupied atomic orbitals on two atoms overlap The two electrons shared in the region of orbital overlap must be of opposite spin Formation of a bond results in a lower potential energy for the system Valence Bond Theory Bond Formation According to valence bond theory atomic orbitals on the central atom are hybridized and form hybrid orbitals The hybrid orbitals on the central atom terminal atoms to form covalent bonds These bonds are called sigma bonds overlap with the atomic orbitals of the In a sigma bond electron density is greatest along the axis of the bond IV Hybridization of Atomic Orbitals Hybrid Orbitals sp hybridization in BeCl2 Orbital diagram of Be 2s 2p Hybridized orbitals sp sp empty p orbitals 7 2 sp 2 hybridization in BF 3 orbital diagram of B 2s 2p Hybridized orbitals sp2 sp2 sp2 empty p orbital 3 a sp 3 hybridization In CH4 Orbital diagram of C 2s 2p Hybridized orbitals sp3 sp3 sp3 sp3 8 b In NH3 Orbital diagram of N 2s 2p Hybridized orbitals sp3 sp3 sp3 sp3 4 sp 3 d hybridization in PCl 5 Orbital diagram of P 3s 3p 3d Hybridized orbitals sp3d empty d orbitals The three hybrid orbitals in the triangular plane are referred to as equatorial while the 9 two hybrid orbitals that are above and below the triangular plane are called axial 5 sp 3 d 2 hybridization in SF 6 Orbital diagram of S 3s 3p 3d Hybridized orbitals sp3d2 empty d orbitals B Summary of orientation of hybrid orbitals V A 1 2 Hybridization in Molecules Containing Multiple Bonds Types of Covalent Bonds Sigma bonds covalent bonds in which the electron density is concentrated between the nuclei