Chapter 11: Theories of Covalent Bonding
Chapter 11: Theories of Covalent Bonding CH 121
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Date Created: 12/01/15
Chapter 11 Lecture Notes Single bonds 0 Single bonds are always 0 bonds because 0 overlap is greater resulting in a stronger bond and more energy lowering Multiple bonds 0 ln multiple bond one of the bonds is a 0 bond and the rest are pi bonds 0 In a molecule like formaldehyde shown at left an spquot2 orbital on carbon overlaps in 0 fashion with the corresponding orbital on the oxygen 0 The unhybridized p orbitals overlap in pi fashion 0 ln triple bonds as in acetylene 2 sp orbitals form a 0 bond between the carbons and 2 pairs of p orbitals in pi fashion to form the two pi bonds 0 The number of electron domains around C determines the hybridization 3 domains spquot2 Pibonds 0 Pi bonds are characterized by 0 Side to side overlap Electron density above and below the internuclear axis Delocalized electrons resonance 0 When writing lewis structures for species like benzene we draw resonance structures to more accurately re ect the structure of the molecule or ion 0 The organic molecule benzene has six 0 bonds around the ring created by overlap of spquot2 orbitals and an unhybridizes p orbital on each carbon atom Resonance o In reality the pi electrons in benzene are not localized but delocalized o The even distribution of the pi electrons in benzene makes the molecule unusually stable Sigma bonds in C2H4 0 One of the C spquot2 hybrid bonds is used to create a sigma bond with C the other 2 form sigma bonds with H Pi bonding in C2H4 o The unused p orbital on each C atom contains an electron and this p orbital overlaps the p orbital on the neighboring atom to form the pi bond Bond order 0 Fractional bond orders occur in molecules with resonance structures Calculating bond order 0 Considering C03 2 0 Bond order total number e pairs used for a type of bondtotal number of bonds of that type CO bond order 4 e pairs in CO bonds3 CO bonds 133 0 Bond length 0 Bond length is the distance between the nuclei of 2 bonded atoms 0 Depends on size of bonded atoms 0 Bond distance measured in Angstrom units where 1 A 10quot10 m o Depends on bond order where the 2 atoms involved in the bond are the same 0 Bond strength 0 Measured by the energy required to break a bond 0 Bond Strength kJmoll CC 346 CC 602 CC 835 o 2 theories of bonding o Valence bond theory linus pauling Valence electrons are localized between atoms or are lone pairs Half lled atomic orbitals overlap to form bonds 0 Molecular orbital theory 0 Robert mullikan Valence electrons are delocalized Valence electrons are in orbitals Called molecule orbitals spread over entire molecule Molecular orbital theory MO 0 Though valence bond theory effectively conveys most observed properties of ions and molecules there are some concepts better represented by molecule orbitals MO can have a maximum of 2e with opposite spin MO has a de nite energy associated with it Can be visualized with a contour representation like atomic orbitals ln MO theory we invoke the wave nature of electrons If waves interact constructively the resulting orbital is lower in energy a bonding molecular orbital If waves interact destructively the resulting orbital is higher in energy an antibonding molecular orbital Schroedinger39s Wave functions 0 O 0 Each Y corresponds to an orbital Region of space within which an electron is found and describes an allowed energy state of an electron Y doesn39t describe the exact location of the electron Yquot2 is proportional to probability of nding an electron at a given point Molecular orbital theory diatomic molecules 0 O 0 Number of MO39s number of atomic orbitals used Bonding MO is lower in energy than atomic orbitals Antibonding MO ls higher Electrons assigned to MO39s of higher and higher energy Molecular orbital theory H2 0 Bonding and antibonding sigma MO39s are formed from ls orbitals on adjacent atoms again we create as many MO39s as we had AO39s MO theory H2 0 ln H2 the 2 electrons go into the bonding molecular orbital The MO bond order is onehalf the difference between the number of bonding and antibonding electron Bond order 12 number bonding electrons number of antibonding electrons For hydrogen H2 with 2 electrons in the bonding MO and none in the antibonding M0 the bond order is c 12 20 2 1 MO theory diatomic molecules 0 Why don39t noble gases exist as diatomic molecules 0 In the case of He2 the bond order would 12 22 0 0 Therefore He2 doesn39t exist 0 If the bond order gt 0 molecule is more stable than separate atoms 0 If bond order 0 molecules is no more stable than separate atoms so will not form 0 MO theory 0 Predicting the existence of the He2 ion 0 The resulting MO diagram looks like this 0 Bond order would be c 12 21 5 o The smaller pblock elements in the 2nd period have a sizeable interaction between the s and p orbitals o This ips the order of the o and pi molecule orbitals in these elements 0 If bond order gt 0 the molecule is more stable than individual atomsions MO theory 0 and pi bonding from p orbitals O For atoms with both 5 and p orbitals there are 2 types of interactions 0 The s and p orbitals that face each other overlap in 0 fashion 0 The other 2 sets of p orbitals overlap in pi fashion 0 There are both 0 and pi bonding molecular orbitals and o and pi antibonding molecular orbitals MO diagrams for 2ndrow diatomic molecules 0 Higher bond order correlates with greater bond strength and shorter bondlength