1. Chapter 9 Covalent Bonding

2. Covalent bond Sharing of electrons –Nonmetal- nonmetal – electronegativity difference less than 1.7

3. Molecule Base unit of a covalent bond

4. Diatomic molecule Naturally formed 2 atom groups –H 2, N 2, 0 2, F 2, Cl 2, Br 2, I 2

5. 2 types of covalent bonds Polar covalent- bonding electrons are shared unequally Nonpolar covalent- bonding electrons are shared equally –Balanced distribution of charge

6. How do you tell the difference in polar and nonpolar? If the 2 atoms have a difference in electronegativity the bond is polar, if no difference then the bond is nonpolar If polar bonds draw the electron to one side of the molecule and the other side has none the molecule is polar

7. Bond length Distance between two bonded atoms at their minimum potential energy –Average distance between bonded atoms

8. Bond energy Energy required to break a bond

9. Coordinate covalent bond One atom donates both electrons to the bond

10. Nomenclature for covalent bonds (nonmetal- nonmetal) 2 systems: –Stock- name(+), (+) oxidation # in roman numerals, name(-) –Classical- use prefixes except for a single (+) ion *Never use mono first*

11. Prefixes Mono Di Tri Tetra Penta Hexa Hepto Octa Nano Deca

12. Naming acids Binary: hydro + root + ic acid Ternary: polyatomic, drop –ate, add –ic acid

13. Oxyacids +1 0 per- ate per- ic Memory -ate - ic -1 0 -ite -ous -2 0 - hypo-ite hypo-ous

14. Lewis structures Use electron dot diagrams to show bonding and electron arrangement

15. Symbols uses in molecular structural formulas unshared pair- (lone pair) pair of electrons that is not involved in bonding- belongs to one atoms single bond- 2e - shared double bond- 4e - shared triple bond- 6e - shared

16. Structural formula Indicates kind, number, arrangement, and bond type in a molecule

17. Sigma bond  - electrons are shared along the bond axis

18. Pi bond  – electrons are shared above and below axis

19. Resonance More than one valid Lewis structure can be drawn

20. VSEPR theory Valance shell electron pair repulsion –Model for molecular geometry –Bond angles –Arrangement minimizes repulsion of e- around the central atom –Molecules adjust their shape, so that valence e- are as far apart as possible

21. 3 types of repulsion 1.Unshared-unshared 2.Unshared- shared 3.Shared-shared

22. hybridization A process in which atomic orbitals are mixed to form new identical hybrid orbitals

23. Hybrid orbitals Orbital of equal energy produced by the combination of two or more atomic orbitals (sp, sp 2, sp 3 )

24. Intramolecular forces Forces within a molecule that hold atoms together

25. Intermolecular forces Forces of attraction between molecules (Van der Waals)

26. Dipole Molecule that has two poles (polar)

27. Dipole moment Measure of the strength of the dipole and is a property that results from the asymmetrical charge distribution in a polar molecule- depends on size and distance  Qd

28. Van der Waals forces Groups of intermolecular forces

29. Dipole- dipole force Between polar molecules

30. Induce dipole A normally nonpolar molecule is transformed into a dipole

31. Hydrogen bonding Intermoleculer (Van der Waals) force in which H bonded to a highly electronegative atom is attracted to an unshared pair of an electronegative atom in a nearby molecule

32. London Dispersion forces (dispersion force) result from the constant motion of e - ’s and the creation of instanteous dipoles –Force generated in a temporary dipole interaction –Most important Van der Waals force –Proposed by Fritz London in 1930 –Strength increases with the number of e- in the interacting atoms