Ch 3 Lecture 1 Simple Bonding Theory

1
Ch 3 Lecture 1 Simple Bonding Theory

• Lewis Structures
• Lewis Bonding Model
• Bonds between atoms consist of one or more shared
  pairs of electrons
• Unshared electrons exist as lone pairs on atoms
• Both bonding pairs and lone pairs influence
• Shape
• Reactivity
• Follows the Octet Rule
• 8 valence electrons (s2p6) is particularly stable
  
• Hydrogen prefers only 2 (1s2) valence electrons
• Examples

2

• Expanded Shells
• Elements in the third or higher shells can have gt
  8 valence electrons
• Empty d-orbitals are used to hold excess
  electrons (s2p6d10 18e- max)
• Examples
• Resonance ability to draw multiple Lewis
  structures for the same molecule
• SO3
• Resonance structures are
• Taken as a set to represent the true structure
• Interconverted by movement of e- only
• Separated by double headed arrows

3

• Isoelectronic resonance structures with
  identical electronic structures
• SO3
• CO32-
• Resonance structures may be electronically
  different as well amides
• More resonance structures means a lower energy
  for the compound. Spreading the electrons over
  more atoms lets them occupy more space.
• Polarity
• Electronegativity ability of an atom to attract
  shared electrons
• Determines polarity of bonds
• Will be discussed more later
• Table 3-1 of your text lists electronegativities
• Fluorine has the highest value

4
Fig 8.2
5

• Polar Bond one between atoms of different
  electronegativities
• Examples
• Use electronegativities to predict /- parts of
  the molecule
• Nonpolar bond one between atoms of the same
  electronegativity
• Formal Charge tool to evaluate resonance
  structures and to explain reactivity
• Apparent charge on an atom based on its Lewis
  structure
• Equation
• An ions charge sum of the formal charges of
  its atoms
• Uses
• Find best resonance structure by minimizing
  charge separation
• Find best resonance structure by matching charge
  / electronegativity

6

• 5. Examples
• Thiocyanate SCN-
• Cyanate OCN-
• Fulminate CNO-
• Ex. 3-1

7

• 6. Expanded Shells can reduce Formal Charge

8

• Be and B Compounds
• BeX2 and BX3 compounds have lt 8 electrons around
  Be/B if single bonds
• Be/BX double bonds create large charge
  separation
• Solids tend to have extended structures relieving
  charge separation/octet
• Monomers are reactive as Lewis Acids (e- pair
  acceptors), even though the small atom size
  allows less than an octet around Be/B

MO calculation
9

• Valence Shell Electron Repulsion Theory (VSEPR)
• Using Lewis structures to predict molecular
  geometry
• Electrons repel each other due to (-) charge
• e- pairs exist due to quantum mechanical rules
  allowing shared orbitals
• Molecules adopt geometries that maximize e- pairs
  separation
• Steric Number (SN) number of atoms and lone
  pairs around a central atom and determines the
  molecules shape
• CO2 SN 2 Shape linear
• SO3 SN 3 Shape trigonal
• Procedure for using VSEPR
• Find the Lewis structure for the molecule
• Determine SN
• Match SN to the appropriate geometry
• Figure 3-8 shows the expected geometries for
  molecules with no lone pairs on the central atom

10
(No Transcript)
11

• 1. Origin of the square antiprism
• SN 5 and SN 7 are not regular because there
  are no regular solids with those numbers of
  corners
• Lone Pair Repulsion
• Lone pairs are as important as atom in VSEPR
• Irregular structures result when the central
• atom has lone pairs