SHAPES OF MOLECULES

REMINDER ABOUT ELECTRONS  Electrons have negative charges  Negative charges “repel” each other  In molecules, electrons want to get as far away from each other as possible  As a result, this repulsion of electrons leads to the shape of the molecule

SHAPE OF MOLECULES  There is a simple model used to determine the shape of molecules.  VSEPR (Valence Shell Electron Pair Repulsion)  A simple model that predicts the general shape of a molecule based on the repulsion between both the bonding and nonbonding electron clouds  Always based on the CENTRAL atom

DEFINITION  Electron cloud: Any type of bond (single, double or triple) or any set of unshared pairs of electrons.  Unshared pair of electron: any pair of electrons not involved in a covalent bond.

2 ELECTRON CLOUDS  Example: CO 2  Count the number of electron clouds surrounding the central atom.. O=C=O · · ··  There are 2 double bonds around the central Carbon (C)  Thus, there are 2 electron clouds

2 ELECTRON CLOUDS  Electrons have a negative charge and repel each other  Thus, the electrons that maintain the bond will move as far away from each other as possible  If the central atom has only two electron clouds, it will have a linear shape

LINEAR  View the clip  NOTE: That the attached atoms are 180° apart from each other

3 ELECTRON CLOUDS  Again count the number of electron clouds around the central atom (SO 3 ).. : O :.. |.. : 0 - S = O ·· ··

3 ELECTRON CLOUDS  Again the electron clouds want to move as far from each other as possible  When the central atom has 3 electron clouds surrounding it, the molecule has a trigonal planar shape

TRIGONAL PLANAR  View the clip  NOTE: The attached atoms are 120° apart from each other

3 ELECTRON CLOUDS  What happens if one of the electron clouds is an unshared pair of electrons  O 3  Do the Lewis dot structure for this in your notes

3 ELECTRON CLOUDS  Notice the following:  You have one double bond  You have one single bond  You have one unshared pair of electrons  When you view a molecule, you can’t see the unshared pair of electrons  This creates a bent shape

BENT  View the clip  NOTE: That the attached atoms and unshared pair are 120° apart from each other  Since you can’t see the unshared pair, the molecule looks bent

4 ELECTRON CLOUDS  Again count the number of electron clouds  CCl 4  Draw the Lewis dot structure in your notes

4 ELECTRON CLOUDS  Same as before, the electron clouds want to get as far from each other as possible  When the central atom has 4 electron clouds surrounding it, you get a tetrahedral

TETRAHEDRAL  View the clip  NOTE: That the attached atoms are 109.5° apart from each other

4 ELECTRON CLOUDS  What happens if you have 3 atoms bound to a central atom with one unshared pair  NH 3  Do the Lewis dot structure for this in your notes

4 ELECTRON CLOUDS  Notice the following:  You have three single bonds  You have one unshared pair of electrons  When you view a molecule, you can’t see the unshared pair of electrons  This creates a pyramidal shape

PYRAMIDAL  View the clip  NOTE: That the attached atoms and unshared pair are 109.5° apart from each other  Since you can’t see the unshared pair, the molecule looks like a pyramid

4 ELECTRON CLOUDS  What happens if you have 2 atoms bound to a central atom with two unshared pairs  H 2 O  Do the Lewis dot structure for this in your notes

4 ELECTRON CLOUDS  Notice the following:  You have two single bonds  You have two unshared pairs of electrons  When you view a molecule, you can’t see the unshared pair of electrons  This creates a bent shape

BENT  View the clip  NOTE: That the attached atoms and unshared pair are 109.5° apart from each other  Since you can’t see the unshared pair, the molecule looks like a pyramid

TRY THESE  HANDOUT  For each of the following, write the Lewis structure and indicate the shape: 1. CBr 4 2. CS 2

POLARITY OF MOLECULES  We’ve already discussed the difference between nonpolar, polar and ionic bonds (electronegativity difference)  Molecular shape is important for determining the polarity of a molecule.  Covalently bonded molecules can be polar or nonpolar based on the shape of the molecule

EXAMPLE  Let’s look at the shapes of  H 2 O  CF 4  First of all:  H-O bond of water has an electronegativity difference of 1.4 (polar covalent)  C-F bond of CF 4 has an electronegativity difference of 1.5 (polar covalent)

EXAMPLE  Since both H 2 O and CF 4 have polar covalent bonds, we would expect both molecules to be polar covalent  This is not the case  IN YOUR NOTES: Draw the molecular shape of both:  H 2 O  CF 4

ANSWER – THINK OF TUG-O-WAR  Water is a bent molecule  As a result water is a polar molecule  You have a partial charge of σ- for O and σ+ for H  CF 4 is a tetrahedral molecule  Because of the shape you have a nonpolar molecule  Even though you have partial charges, the charges cancel out because of the shape

WRITE WHICH SHAPES ARE POLAR/NONPOLAR?  Take a look at your handout that shows the shape of different molecules.  Which shapes do you think are polar?  Which shapes do you think are nonpolar?

ANSWER  Polar shapes (have lone pairs):  Bent  Pyramidal  Nonpolar shapes (do not have lone pairs) :  Linear  Trigonal planar  Tetrahedral

TRY THE FOLLOWING  Determine the Lewis structure of the following. Are the molecules polar or nonpolar? 1. Cl 2 O 2. CO 2 3. NF 3

ANSWER 1. Bent, polar molecule 2. Linear, nonpolar molecule 3. Pyramidal, polar molecule