Molecular shapes Balls and sticks

Learning objectives  Apply VSEPR to predict electronic geometry and shapes of simple molecules  Distinguish between polar and nonpolar bonds in molecules  Predict polarity of simple molecules from bond polarity and molecular shape

Roadmap to polarity  Establish skeleton of molecule  Determine Lewis dot structure using S = N – A  Determine electronic geometry using VSEPR  Identify molecular geometry from molecular  Count number of polar bonds  Perform polarity analysis

Valence shell electron pair repulsion  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 molecular shape  First approach to this problem is to consider repulsion between groups of electrons (charge clouds)

Electron groups (clouds) minimize potential energy  Valence shell electron pair repulsion (VSEPR)  Identify all groups of charge: non- bonding pairs or bonds (multiples count as one)  Bonded atoms – single, double or triple count as 1  Lone pairs count as 1  Distribute them about central atom to minimize potential energy (maximum separation)

Choices are limited  Groups of charge range from 2 – 6  Only one electronic geometry in each case  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)

Total number of groups dictates electronic geometry  Octet rule:  Two – linear  Three – trigonal planar  Four – tetrahedral  Additional possibilities (expand octet):  Five – trigonal bipyramidal  Six - octahedral

Stage 3: Molecular shape:  What you get from electronic geometry considering atoms only  Same tetrahedral electronic geometry – different molecular shape

Two groups: linear  Except for BeH 2, all cases with two groups involve multiple bonds

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

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  Note: 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°

Representations of the tetrahedron

Groups of charge Lone electron pairs Electronic geometry Molecular shape 20Linear 30Trigonal planar 31 Bent 40Tetrahedral 41 Trigonal pyramid 42TetrahedralBent

Important properties related to polarity  Solubility: polar molecules dissolve in polar solvents; nonpolar molecules dissolve in nonpolar solvents  Oil (nonpolar) and water (polar) don’t mix  Ammonia (polar) dissolves in water  Melting and boiling points  Polar substances have high intermolecular forces:  Melting and boiling points are much higher than with nonpolar substances (H 2 O is a liquid, CO 2 is a gas)

Roadmap to polarity  Establish skeleton of molecule  Determine Lewis dot structure using S = N – A  Determine electronic geometry using VSEPR  Identify molecular geometry from molecular  Count number of polar bonds  Perform polarity analysis

Polar bonds and polar molecules  Not all molecules containing polar bonds will themselves be polar.  Need to examine the molecular shape  Ask the question:  Do the individual bond polarities cancel out?  If so, non polar. If not, polar.

Consider some examples  In CO 2 (linear molecule) the two polar bonds oppose each other exactly  In chemical tug-o-war there is stalemate

The most important polar molecule  In BF 3 the three bonds cancel out – tug of war stalemate  In H 2 O (bent) the polar bonds do not directly oppose – no stalemate  Lone pair also adds some component  Overall net polarity  Consequence of polarity: H 2 O is a liquid, CO 2 is a gas

Symmetry and polarity  If the molecule “looks” symmetrical it will be nonpolar  If the molecule “looks” non-symmetrical it will be polar

Rules of thumb for evaluation of polarity  Presence of one lone pair of electrons  Only one polar bond  Always polar molecules  Two or more polar bonds  Do polar bonds perfectly oppose?  If no, polar molecule

Two bonds  Equal bonds oppose (linear)  Nonpolar (CO 2 )  Unequal bonds oppose (linear)  Polar (HCN)  Equal bonds do not oppose (bent)  Polar (H 2 O)

Three bonds  Equal bonds oppose in trigonal planar arrangement  Nonpolar  Unequal bonds in trigonal planar arrangement  Polar