Chapter 8 Section 3

1. Chapter 8 Section 3 Molecular Structures

2. Structural Formulas • Molecular formulas tell you the type and number of each atom in a molecule. • Example: C2H5OH • There a different models that can be used to represent a molecule.

3. Structural Formulas • One of the most useful is the structural formula, which uses letter symbols and bonds to show the relative positions of atoms. • You can predict the structural formula of a molecule by drawing the Lewis structure of it.

4. Rules for drawing Lewis Structures • 1. Predict the central atom – usually the atom in the molecule closest to the left side of the periodic table. • 2. Determine the number of electrons available for bonding – This is equal to the sum of the valence electrons for all atoms present. • 3. Bond each atom – use two electrons to form a single bond between each atom present. • 4. Distribute remaining electrons.

5. Rules for drawing Lewis Structures • 5. Check that the octet rule is satisfied for all atoms (except hydrogen) • If not, begin making double and triple bonds until the octet rule is obeyed by all atoms (except hydrogen).

6. Lewis Structure Example • NH3 • Predict central atom – N • Sum electrons – 5 + 1 + 1 + 1 = 8 • Bond each atom using a pair of electrons • Distribute remaining electrons • Check octet rule

7. Lewis Structure Example • You try…H2O

8. Lewis Structure Example • CO2 • Predict central atom – C • Sum up electrons – 4 + 6 + 6 = 16 • Bond each atom • Distribute remaining electrons • Check for octet rule…is it satisfied?

9. Lewis Structure example • CO2 continued • The octet rule was not satisfied for each element, so we now start over using a double bond instead of a single bond and repeat the process

10. Lewis Structure Example • You try CO

11. Lewis Structures for Polyatomic Ions • Only difference is that these molecules have charges. • If the charge is negative, add that number to the total number of valence electrons. • If the charge is positive, subtract that number from the total number of valence electrons.

12. Polyatomic Example • PO4-3

13. Polyatomic Example • You try NH4+

14. Exceptions to the Octet Rule • Suboctets: A few elements can form bonds and have less than an octet. An example is Boron • BH3 • Expanded Octets: Usually can occur in elements found in periods 3 or higher because of its d energy levels. These have more than an octet. Example: Phosphorous • PCl5

15. Chapter 8 Section 4 Molecular Shapes

16. VSEPR Model • The shape of a molecule determines many of its physical and chemical properties. • The molecular geometry or shape of a molecule can be determined once a Lewis structure is drawn. • The used to determine the molecular shape is referred to as the Valence Shell Electron Pair Repulsion Model or VSEPR Model

17. VSEPR Model • This model is based on the arrangement that minimizes the repulsion of shared and unshared electron pairs around the central atom. • The electron pairs in a molecule repel each other and this force causes the atoms in the molecule to be positioned at fixed angles relative to one another. • The angle formed by two terminal atoms and the central atom is the bond angle.

18. Molecular Shapes

19. Molecular Shapes Linear Example: BeCl2 Trigonal Planar Example: AlCl3

20. Molecular Shapes Tetrahedral Example: CH4 Trigonal Pyramidal Example: PH3

21. Molecular Shapes Bent Example: H2O TrigonalBipyramidal Example: NbBr3

22. Molecular Shapes Octahedral Example: SF6