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Molecular shapes A simple matter of balls and sticks.

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Presentation on theme: "Molecular shapes A simple matter of balls and sticks."— Presentation transcript:

1 Molecular shapes A simple matter of balls and sticks

2 Learning objectives  Describe underlying principles that govern theories of molecular shapes  Use Lewis dot diagrams to predict shapes of molecules using VSEPR

3 Valence shell electron pair repulsion  In order to understand properties like polarity, we need to predict molecular shapes  Lewis dot structure provides 2D sketch of the distribution of the valence electrons among bonds between atoms and lone pairs; it provides no information about the shape of the molecule

4 A hierarchy of models  VSEPR  Consider the problem in terms of electrostatic repulsion between groups of electrons (charge clouds, domains)  Valence bond theory  Acknowledges the role of orbitals in covalent bonding  Molecular orbital (MO) theory (the “real” thing)  Accommodates delocalization of electrons - explains optical and magnetic properties

5 Electron groups (clouds) minimize potential energy  Valence shell electron pair repulsion (VSEPR)  Identify all of the groups of charge: non-bonding pairs and bonds (multiples count as one)  Distribute them about the central atom to minimize potential energy (maximum separation of the groups)  This specifies the electronic geometry (also known as electron domain geometry or sometimes confusingly as molecular geometry)

6 Choices are limited  Groups (domains) of charge range from 2 – 6  Only one electronic geometry in each case  However, more than one molecular shape follows from electronic geometry depending on number of lone pairs  One surprise: the lone pairs occupy more space than the bonded atoms (with very few exceptions)  Manifested in bond angles (examples follow)  Molecular shape selection (particularly in trigonal bipyramid)

7 Two groups: linear  Except for BeH 2 (Be violates octet rule), all cases with two groups involve multiple bonds

8 Three groups: trigonal planar  Two possibilities for central atoms with complete octets:  Trigonal planar (H 2 CO)  Bent (SO 2 )  BCl 3 provides example of trigonal planar with three single bonds  B is satisfied with 6 electrons – violates octet rule

9 Four groups: tetrahedral  Three possibilities:  No lone pairs (CH 4 ) - tetrahedral  One lone pair (NH 3 ) – trigonal pyramid  Two lone pairs (H 2 O) – bent  Lone pairs need space: H-N-H angle 107°H-N-H angle 107° H-O-H angle 104.5°H-O-H angle 104.5° Tetrahedral angle 109.5°Tetrahedral angle 109.5°

10 Representations of the tetrahedron

11 Five groups of charge: trigonal bipyramid – most variations  Two different positions:  Three equatorial  Two axial  Equatorial positions are lower energy:  Lone pairs require occupy these locations preferentially

12 Five bonds, no lone pairs

13 Four bonds, one lone pair  Lone pair dictates geometry: equatorial position has lower energy than axial

14 Three bonds, two lone pairs  Both lone pairs occupy equatorial positions – lower energy than in axial

15 Two bonds, three lone pairs  The trend continues: all equatorial positions filled – lowest energy

16 Octahedron has six identical positions and high symmetry

17 No lone pairs  High symmetry

18 One lone pair  All positions are equally probable  Symmetry reduced

19 Two lone pairs  Minimum energy has axial symmetry, lone pairs lie along straight line



22 Molecules with multiple centers  A central atom is any atom with more than one atom bonded to it  Perform exercise individually for each atom  Electronic geometry and molecular shape will refer only to the atoms/lone pairs immediately attached to that atom

23 Taking it to the next level: acknowledging orbitals  VSEPR is quite successful in predicting molecular shapes based on the simplistic Lewis dot approach  But our understanding of the atom has the electrons occupying atomic orbitals  How do we reconcile the observed shapes of molecules with the atomic orbital picture of atoms

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