1. Molecules

2. Objectives Write the electron dot structure for an atom. Explain how covalent bonds form molecules.

3. Electron Dot Structures Dots can be used to represent v/e. Write electron dot structures for K, P, S, and Br.

4. Covalent Bonds How do non-metal atoms bond together? Example: chlorine gas, Cl 2 2e- Both nuclei are attracted to the molecular orbital. Both nuclei repel each other. This equilibrium establishes the bond length. + + covalent bond: an attraction between non-metal atoms sharing v/e in a molecular orbital

5. Molecules molecule: group of neutral atoms held together with covalent bonds molecular formula: indicates exact number and kinds of atoms in the molecule IMPORTANT ionic compound formula: ratio of the ions in the crystal CaCl 2 (1:2 ratio) sucrose: C 6 H 12 O 6

6. Objective Be able to draw a Lewis diagram when given a molecular formula.

7. Lewis Structures Draw Lewis structures for molecules containing the following elements: nitrogen and fluorine sulfur and chlorine hydrogen and oxygen (water)

8. Lewis Structures Double and triple covalent bonds can also form: H 2 CO (double) ClCN (triple) Stick Diagrams C 6 H 14

9. Objectives Understand the concept of VSEPR theory. Use VSEPR theory to determine the shape(s) present in a molecule. Be able to draw Lewis diagrams for various shapes.

10. VSEPR Theory Valence Shell Electron Pair Repulsion Theory: each pair of electrons (bonding pair or unshared pair) will repel; molecule will adjust shape to maximize the angles between each pair

11. Molecular Shapes CH 4 methane tetrahedral (4 single) NH 3 ammonia pyramidal (3 single) H 2 S dihydrogen sulfide bent (2 single)

12. Molecular Shapes trignonal planar (2 single, 1 double) H 2 CO formaldehdye CO 2 carbon dioxide linear (2 double) HCN hydrogen cyanide linear (1 single, 1 triple)

13. Objectives Be able to determine the polarity of a covalent bond using a table of electronegativities. Be able to determine the polarity of a molecule based on the shape(s) present and the polarity of the covalent bonds within the molecule. Understand and apply the concept of molecular symmetry.

14. Bond Polarity Electrons are not always shared evenly. H F Atom with higher electronegativity attracts the e- pair more! -- ++ polar bond: unevenly shared covalent bond Br non-polar bond: evenly shared covalent bond 2.2 4.0 3.0

15. Bond Polarity What is the polarity of a H – C bond? H – O? K – Cl? 0.5 – 1.9 = polar covalent (uneven sharing) Electronegativity difference determines bond polarity: 0.0 – 0.4 = non-polar covalent (even sharing) 2.0 – above = ionic bond (e- transfer)

16. Molecular Polarity Bond polarity and molecular shape must be considered when determining molecular polarity. dipole: a molecule that has an uneven distribution of charge (  + and  - sides can be separated with a line) ++ -- -- ++ NH 3 H2OH2O

17. Molecular Polarity A molecule with all non-polar bonds is called non-polar molecule H-C bonds are always non-polar (example: hydrocarbons) Due to symmetry, a molecule with polar bonds can be a non-polar molecule.  

18. Objectives Be able to determine the type of intermolecular bonding present in a molecular compound. Be able to predict the state of a molecular compound based on the type of intermolecular bonding present and the mass of the molecules.

19. Intermolecular Bonds intermolecular bond: a bond between molecules The state of a substance is determined by the strength of these bonds. Non-polar molecules tend to be gases (don’t attract): O 2, N 2, CO 2

20. London Dispersion Forces London dispersion force: caused by random distributions of electrons; brief partial charges cause attraction; more electrons = stronger bonds non-polar molecule gas: < 70 e- liquid: 70 e- to 100 e- solid: > 100 e-      

21. Dipole Interactions dipole interaction: dipoles attract each other; liquid or solid London force determines state: liquid: < 100 e- solid: > 100 e-     

22. Hydrogen Bonds hydrogen bond: strong dipole interaction; occurs between molecules with –OH or –NH always liquid or solid (depends on strength of London forces)

24. Hydrogen Bonds H-bonds occur in both H 2 O and DNA.

25. Objectives Be able to name molecular compounds. Know the diatomic elements. Recognize elements that produce macromolecules.

26. Molecular Names Covalent Compound or Molecular Compound Nomenclature Prefixes 1mono-(mon-) 2di- 3tri- 4tetra-(tetr-) 5penta- (pent-) 6hexa-(hex-) 7hepta-(hept-) 8octa-(oct-) 9nona-(non-) 10deca-(dec-) use prefixes that match the number of atoms in the formula end the second name with “-ide” never start the first name with “mono-“; only use it for the second name The o or a at the end of a prefix is usually dropped when it precedes oxide.

27. Diatomic Elements Many elements exist as paired atom molecules H 2, N 2, O 2, and halogens

28. Macromolecules macromolecule: a huge molecule C and Si produce macromolecules because each element can form four bonds per atom diamond (C) quartz (SiO 2 )

30. Properties of Covalent Compounds Composed on non-metals only Covalent bonding—exist as molecules. Solid, liquid, or gas Usually non-conductors of heat and electricity

31. Polyatomic Ions and Coordinate Covalent Bonds (extra-credit) Polyatomic ions are essentially charged molecules (usually with additional electrons). Draw Lewis structures for NO 2 – and OH – coordinate covalent bond: a bond that forms when one atom donates both electrons to a bond Examples: CO, SO 3, CO 3 2–