Laboratory 06 MOLECULAR GEOMETRY AND POLARITY

1. 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 Boiling point, Freezing point, Chemical reactivity Polar Non Polar Reactivity Toxicity

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

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. Two Rules Construction of Lewis Structures 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 Bent Trigonal planar

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. 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 Bond Polarity m

17. Polarity • Covalent bonds and molecules are either polar or nonpolar • Polar: • 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.