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 number and kinds of atoms in a molecule H2O2H2O2 O O H H hydrogen peroxide

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)

9. Objective Understand the concept of VSEPR theory. Use VSEPR theory to determine the shape(s) of a molecule. Be able to identify common molecular 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 between atoms. 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? A bond between H – O? A bond between 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 the polarity of a molecule. dipole: a polar molecule with polar bonds in which partial charges can be separated ++ -- -- ++ NH 3 H2OH2O

17. Molecular Polarity A molecule with all non-polar bonds is called a non-polar molecule. Hydrocarbons (with just H and C) are always non-polar molecules. 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 Covalent bonds form within molecules (polar or non-polar). intermolecular bond: a bond between molecules State of substance determined by strength Non-polar molecules tend to be gaseous because they don’t attract each other. Examples: O 2, N 2, CO 2

20. London Dispersion Forces London dispersion force: an intermolecular force caused by random distributions of electrons Brief partial charges cause attraction—more electrons causes stronger bonding! Generally, in a non-polar molecule… gas: < 70 e- liquid:70 e- to 100 e- solid: > 100 e-      

21. Dipole Interactions dipole interaction: an attraction between dipoles that always results in a liquid or solid London force strength will determine state: liquid: < 100 e- solid: > 100 e-     

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

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

24. Objectives Be able to draw a Lewis diagram for a polyatomic ion. Be able to draw a Lewis diagram for a molecule or ion containing coordinate covalent bonds.

25. 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-“ The o or a at the end of a prefix is often dropped when the word following the prefix begins with another vowel; however, it is okay to have an a before the i in iodide

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

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

29. 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

30. 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–