1. MOLECULAR GEOMETRY AND POLARITY
Laboratory 06
MOLECULAR GEOMETRY AND POLARITY

2. The three dimensional (3D) structure of a molecule
-How its atoms are connected and arranged in space
Molecular geometry
Polar
Boiling point, Freezing point, Chemical reactivity
Non Polar
Reactivity
Toxicity

3. Determination of the correct 3D structure of a
molecule
Strategy
Electron domain geometry
Lewis structure
Molecular geometry

4. Background- Lewis structure
Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule

5. Construction of Lewis Structures
Two Rules
Total # of valence electrons – the total number of valence electrons must be accounted for, no extras, none missing.
Octet Rule – every atom should have an octet (8) electrons associated with it.
Hydrogen should only have 2 (a duet).

6. Drawing Lewis Dot Structures
Determine the total number of valence electrons.
Determine which atom is the “central” atom.
Stick everything to the central atom using a single bond.
Fill the octet of every atom by adding dots.
Verify the total number of valence electrons in the structure.
Add or subtract electrons to the structure by making/breaking bonds to get the correct # of valence electrons.
Check the “formal charge” of each atom.

7. Formal Charge of an Atom
Formal charge = number of valence electrons – number of
bonds – number of non-bonding electrons.

8. Determination of the electron domain geometry around the central atom
A) Lewis structures do not indicate the three dimensional shape of a molecule. They do not show the arrangement space of the atoms, what we call the molecular geometry or molecular structure.
B) Molecules have definite shapes and the shape of a molecule controls some of its chemical and physical properties.
Valence Shell Electron Pair Repulsion Theory - VSEPR - predicts the shapes of a number of molecules and polyatomic ions.

9. VSEPR Theory Predicts the molecular shape of a bonded molecule
Electrons around the central atom arrange themselves as far apart from each other as possible
Unshared pairs of electrons (lone pairs) on the central atom repel the most
lone pairs of electrons are spread out more broadly than bond pairs.
(repulsions are greatest between two lone pairs, intermediate between a lone pair and a bond pair, and weakest between two bonding pairs of electrons)
Repulsive forces decrease rapidly with increasing interpair angle
(greatest at 90o, much weaker at 120o, and very weak at 180o)

10. Electron domain geometry
-Describe the geometric arrangement around the atom
-Compose, Single lone pair of electrons or chemical bond; a single, double, or triple bond
-There are five basic electron domain geometries possible.

11. Molecules with no lone pairs
Trigonal planar
Trigonal bipyramidal
Trigonal planar
Trigonal bipyramidal

12. Molecules with 3 electron groups
Trigonal planar
Trigonal planar
Trigonal planar
Bent

13. Molecules with 4 electron groups
Trigonal pyramidal
Bent

14. Molecules with 5 electron groups
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal

15. Molecules with 6 electron groups

16. Bond Polarity
Due to differences in electronegativities of the bonding atoms
If Den = 0, bond is nonpolar covalent
If 0 < Den < 2, bond is polar covalent
If Den > 2, bond is ionic m

17. Polarity Covalent bonds and molecules are either polar or nonpolar
Electrons unequally shared
More attracted to one nuclei
Nonpolar:
Electrons equally shared
Measure of polarity: dipole moment (μ)

18. Molecular Polarity Overall electron distribution within a molecule
Depends on bond polarity and molecular geometry
Vector sum of the bond dipole moments
Lone pairs of electrons contribute to the dipole moment
Consider both magnitude and direction of individual bond dipole moments
Symmetrical molecules with polar bonds = nonpolar

19. Resonance structures
In chemistry, resonance is a way of describing delocalized electrons within certain molecules or polyatomic ions, where the bonding cannot be expressed by one single Lewis formula.
A molecule or ion with such delocalized electrons is represented by several contributing structures also called resonance structures.

20. Resonance structures
1. Do not violate the octet rule!!! (DO NOT HAVE 5 BONDS TO CARBON!!!)
2.The overall charge of a molecule should not change—atoms may have charges, but the net charge of the entire molecule should not change
3. Place resonance structures inside brackets ([ ]) and use to separate each structure
4. Do not break σ bonds! (e.g. do not break C-C, C-H, C-O, or C-N single bonds)
5.Charges will be preferentially located on atoms of compatible electronegativity. For example, oxygen is more electronegative than carbon; therefore, a negative charge will preferentially be placed on oxygen rather than carbon in the dominant resonance structure.
6. Unless starting with a radical, move electrons in pairs, using a double-headed arrow.
7. Do not “jump” lone pairs from one atom to another. Lone pairs can become π bonds or π bonds can become lone pairs. π Bonds can migrate from one side of a carbon atom to
another.