Chemical Bonds: The Formation of Compounds from Atoms

1. Chemical Bonds: The Formation of Compounds from Atoms Dr. Bixler-Zalesinsky

2. 11.1 Periodic Trends in Atomic Properties

3. Metals and Nonmetals (review)

4. Ionization energy is the energy required to remove an electron; corresponds to their charge

5. Atomic Radii increase going down a group and decrease across a period

6. Valence Electrons Lewis Dot Diagrams

7. Valence Electrons & Per. Table

8. Lewis Structures of an atom shows the valence electrons (ones involved in bonding)

9. Octet Rule • Every atom aspires to have eight electrons in its outermost shell (2 s electrons and 6 p electrons just like the noble gases) • They must borrow (covalent molecules), release or accept (ionic compounds) electrons to get to the eight.

10. Types of bonding

11. Ions and Bonding

12. Ionic bonding occurs between metals (cations) and nonmetals (anions)

13. The nonmetal accepts the electron(s) and the metal donates the electron(s) ionic bond is the attractive between oppositely charged ions

14. Form large crystals; our formulas are the smallest whole number ratios not the true number of atoms

15. The charges must cancel out and equal zero to form stable compounds If you have a +2 ion then you need either two -1 ions or one -2 ion

16. Link to Video clip on Ionic Bonding (1:39)

17. Bonding Covalent v. ionic

18. 11.5 Covalent Bonding: Sharing Electrons • Covalent bonding occurs between two nonmetal atoms • Electrons are shared between two atoms

19. 11.7 Lewis Structures of Molecules • Find the number of valence electrons for each element in the structure • Multiply the number of valence electrons times the number of atoms you have of that element • Determine which element can make the most bonds and put it in the center and attach the other elements to it • Make each atom have 8 valence electrons around it. • Add up the number of electrons you used in the structure. This number must match the total number of electrons you started with

20. Number of Valence Electrons: Page 151 in textbook 1 2 3 4 5 6 7 8 He H 2.1 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Li 1.0 Na 1.5 Ne Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar K 08 Ca 1.0 Kr Br 2.8 Se 2.4 Xe Rb 0.8 Sr 1.0 I 2.5 3 4 3 2 1 0 (number of bonds each can make) S S/D/T S/T S/D S 0 (types of bonds s=single, D= double, T= triple) Cs 0.7 Ba 0.9 Rn Fr 0.7 Ra 0.9

21. Ex. Write the Lewis Dot Diagrams for the following molecules • I2 • H2O • FCl • CF4 • NBr3

22. Molecule (Covalent) Nomenclature • Naming: These binary inorganic molecules are named using prefixes like mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-, • The first element only gets a prefix if there is more than one of it; otherwise, the element name remains the same. • The second element ALWAYS gets a prefix and the ending changes to –ide. • Ex. CO is carbon monoxide (two words not capitalized)

23. HW p. 164 # 2 a – f Write the question and answer!

24. Multiple Bonds Double and Triple Bonds Knowing when NOT to use them is as important as understanding when to use them!

25. Multiple Bonds • Some times using the correct number of electrons will not give you a full octet. When this happens: • 1st double check your math and counting • 2nd see if the atoms involved can make a double or triple bond

26. Double Bonds • O, S, and C can make double bonds with each other than themselves but no others! • Double bond is 4 electrons in a bond • Symbolized by an = sign • Take a look at CO2

27. Triple Bonds • P, N, and C can make a triple bond with each other or themselves • A triple bond is 6 electrons in a bond • The symbol for a triple bond is = • Let’s try N2

28. Molecular Geometry VSEPR Theory and Application

29. Structural Formula • Shows how elements of a molecule are connected to each other

30. VSEPR • V = valence • S = shell • E = electron • P = pair • R = repulsion • Electrons will arrange themselves as far apart from one another as possible • Unbonded pairs take up more room than bonded ones

31. 3-D Hybridized orbitals, shapes, and decision tree

32. Linear Shape

33. Bent Shape

34. Trigonal planar

35. Pyramidal

36. Tetrahedral

37. Video Review (3:21) VSEPR

38. sp hybridization

39. Path to hybridization

40. Sp3 hybridization

41. Video Review (1:36) Hybridization

42. VSEPR Theory of Molecular Geometry

43. VSEPR Theory of Molecular Geometry

44. Shape Decision Tree

45. Polar Covalent v. Nonpolar Covalent

46. Polar and Nonpolar Covalent Bonds • If they are shared equally they are said to be nonpolar bonds if they are not equally shared they are said to be polar bonds • Sharing of electrons has to do with the pull of one element compared to the other element sharing the electron pair. This pulling is called electronegativity (eneg) which increases across the period and up the group • The larger the electronegativity the greater the time the electrons spend with the more electronegative atom giving it a slightly positive charge and because of this imbalance we call this a polar molecule • If the eneg difference lies between 0.5 to 1.6 it is a polar bond

47. Polar or Nonpolar Molecules • Determine the shape of the molecule • Determine how many polar bonds there are in the molecule • If there are NO polar bonds the molecule must be NONpolar. • If there is exactly one polar bond, the molecule is polar. • If there is more than one polar bond, follow the chart below.

48. Molecules with more than one polar bond (assuming polarity is equal)

49. VSEPR Theory of Molecular Geometry

50. Polarity of Molecules with more than one polar bond