2. Covalent bonding The sharing of one or more pairs of elec- trons so that the electron con- figuration fulfills the octet rule.

3. Many atoms will have unshared or lone pairs of electrons in a molecule. Hydrogen does not have any unshared pairs of electrons, because it’s orbital is full with 2 electrons. It always forms only one bond and therefore it cannot be a central atom is a molecule. O-H-O

4. Covalent bonding in water.

6. Representations of Molecules. We will most often be drawing structural formulas.

7. Carbon bonding Carbon has an electron conf. of 1s 2 2s 2 2p 2. This means that this s electrons are paired and the p electrons are not paired. If carbon retains this e- structure, it would form only 2 or 3 bonds rather than 4 as you saw in methane on the previous slide. One of the 2s electrons will jump to the 3 rd 2p orbital by absorbing a little energy. Then carbon has 4 unpaired electrons to form 4 single covalent bonds.

8. Structural formulas of several compounds w/ single bonds. Some models show unshared or lone pairs of electrons whereas some do not.

9. What are the formulas for these molecules? Why does BF 3 not have an octet around the boron?

10. Models showing all electrons. These are not your typical models, but they do show all of the electrons in the energy levels of the bonded atoms.

11. Comparison of ionic and covalent bonds.

12. Single Covalent bonds In a single covalent bond, 2 electrons are being shared equally be two different nuclei. The orbitals of the two atoms overlap.

13. The electron orbitals are shown in a covalent bond.

14. Water molecules and sharing of electrons. Water molecules are polar. That means that they have a slightly positive end and a slightly negative end. Why does this occur?

15. Double and triple bonds Double covalent bonds occur when 2 pair or 4 electrons are shared. Triple covalent bonds occur when 3 pair or 6 electrons are shared.

16. Double and triple covalent bonds Double and triple bonds occur when there are fewer electrons available to share.

17. Coordinate Covalent Bonds These are also called dative bonds in which both electrons that are being shared come from one of the two atoms and fill an empty orbital. These types of bonds are no different than any other. Example is the ammonium ion, NH 4 + Table 16.2, p. 445 – Some Common Covalent Compounds

18. How do you know what shape to draw? Typically the first atom listed is the central atom or the base chain as in hydrocarbons. Ex.: PCl 3 or C 4 H 10 Most everything else is bonded off of the central atom.

19. Steps in drawing a molecule 1. Count and add up the number of valence electrons for each molecule. 2. Test different drawings until each atom has a full octet, except for hydrogen. As the number of pairs of electrons decreases, the number of double or triple bonds would increase. If you give each atom 8 e - and only single bonds and you are one pair short of the valence number you need one double bond. If you are two pairs of e - short, you need a triple bond or two double bonds.

20. A rule of thumb is HONC: hydrogen forms one bond, oxygen tends to form 2 bonds with 2 lone pair of electrons, nitrogen tends to form 3 bonds with one lone pair of electrons and carbon forms 4 bonds Ethane

21. Practice CCl 4 H 2 S H 2 O 2 PCl 5 XeCl 4

22. Exceptions to the octet rule Hydrogen – only forms one single covalent bond Beryllium – only forms two single covalent bonds Boron – only forms three single covalent bonds The two directly above are so small that they cannot lose half or more of their electrons so they share the few valence electrons that they have. Phosphorus can form 5 bonds (plus others) Sulfur can form 6 bonds (plus others such as xenon) Expanded octets occur only around the central atom

23. More Practice SO 2 SO 3 SO 3 2- CO 3 2-

24. Resonance Structures Any molecule in which there is more than one possible Lewis structure that can be drawn. (The molecule itself isn’t moving, but the bonds are in different positions.) Example: NO 2 - Each molecule has double arrow between them and if it carries a charge, brackets around the molecule are included. Work problems, p. 451 #7-12

25. Shape and polarity determine many properties of a chemical such as solubility, melting and boiling points, state of matter and reactivity. We will be using molecular geometry / shape to determine if the molecule is polar or nonpolar.

26. Shapes of simple molecules and ions – VSEPR Theory Valence shell electron pair repulsion Pairs of electrons arrange themselves around the central atom so that they are as far apart from each other as possible. There will be greater repulsion between non-bonded pairs of electrons than bonded pairs. (They will take up more space. Double and triple bonds count as one pair of electrons in VSEPR theory. By loose definition any region of electrons is called a negative charge center.

27. Shapes and bond angles 2 negative charge centers – linear with a bond angle of 180 o BeF 2 CO 2 C 2 H 2 HCN

28. Shapes and bond angles 3 negative charge centers - two possible shapes Trigonal planar – 3 bonding pairs or electrons with a bond angle of 120 o BF 3 C 2 H 4 CO 3 2-

29. Shapes and bond angles 3 negative charge centers - two possible shapes Bent or V-shaped made of 2 bonding pairs and one non-bonded pair of 120 o (since there are only three areas of charge, the non-bonded pair doesn’t make a difference in bond angle) SO 2 NO 2 -

30. Shapes and bond angles 4 negative charge centers – tetrahedral – 4 bonding pairs with a bond angle of 109.5 o CH 4 CCl 4 NH 4 + BF 4 -

31. Shapes and bond angles 4 negative charge centers – tetrahedral geometry– 3 bonding pairs and one non- bonding pair of electrons-trigonal pyramid with a bond angle of 107.3 o Ammonia, NH 3

32. Shapes and bond angles 4 negative charge centers – tetrahedral geometry– 2 bonding pairs and two non- bonding pair of electrons-angular shape with a bond angle of 104.5 o H 2 O

33. 3 basic tetrahedral shapes

34. Shapes and bond angles 5 negative charge centers - trigonal bipyramidal – 90 o, 120 o and 180 o PCl 5 There are several other shapes, but IB doesn’t emphasize.

35. Shapes and bond angles 6 negative charge centers – 90 o and 180 o octahedral square planar – 4 bonding pairs with 2 non-bonding as far apart as possible above and below the plane of the molecule Distorted tetrahedral (see-saw) - 5 bonding pairs with 1 non-bonding

37. Other forms of the basic shapes

38. A review of the shapes

41. Work problems, p. 470-473 #28, 32, 33, 37, 48, 63

42. Polar bonds and molecules A value used to look at polarity is called electronegativity, the higher this value is the greater the attraction for electrons in a shared pair. We will study it further in Ch. 14. The electronegativity difference between two atoms tells you which end has a greater attraction, possibly creating a polar molecule. Table 16.4, p. 462 defines the types of bonds.

43. Polar bond in a molecule.

44. Nonpolar covalent bonds form when electrons are shared equally. Example would include the diatomic molecules or electronegativity differences less than 0.4. Ionic compounds have very large electronegativity differences, greater than or equal to 2.0.

45. Polar molecules Polar molecules fall in-between nonpolar and ionic substances. In a linear molecule, if the electronegativity is between 0.4 – 2.0 it is polar. If the molecule isn’t linear, polarity depends on the shape. If it is the same all the way around and not bent, it is most likely nonpolar. If one end is different than the other, it will be polar. Work problems, p. 466 #21-26, 51, 52

46. Intermolecular Forces Many molecular compounds are liquids and gases, with those that are solids having a very high molar mass. The attractions in-between (intermolecular force) molecules determines the state and melting/boiling point of a substance.

47. 3 states of matter and intermolecular forces.

48. Gas, liquid and solid molecules

49. 3 types of intermolecular forces/attractions from weakest to strongest van der Waals forces – also called London dispersion forces. (weakest) Dipole – Dipole forces H-Bonding – not an actual bond (strongest)

50. van der Waal’s forces Weak intermolecular forces that occur when electrons happen to collect on one end of a molecule, which induces charge on adjacent molecules.

51. van der Waal’s forces Called temporary dipole moments.

52. van der Waal’s forces

53. Dipole moments are temporary, but they cause changes in adjacent molecules, thus creating a very weak intermolecular attraction. Only intermolecular force in non-polar molecules.

54. van der Waal’s forces

55. Dipole-Dipole forces In polar molecules the force of attraction is a more constant force. This means that polar substances tend to have greater attractions and higher melting/boiling points.

56. Dipole-Dipole forces

57. Molecules line up so that the positive and negative ends attract.

58. Dipole-Dipole forces

59. H-bonding Hydrogen bonding is not a true bond, but it is a very strong attraction between molecules. It occurs in molecules where hydrogen is bonded to either nitrogen, oxygen or fluorine. These three elements have a very high electro-negativity, because of their small size and wanting to gain electrons. The attraction between the hydrogen and other element is so high that the hydrogen almost completely loses its electron, becoming nearly 1+ and the N, O, or F attracts the electrons to become nearly a 1-.

61. H-bonding in water & ice

62. Boiling points with h-bonding

63. Crystal lattice or network covalent solid Molecular compounds that are not small, but very large and every attraction is a covalent bond between atoms. They are insoluble solids that are hard with a very high melting points. Examples are diamond, graphite and silicon dioxide (sand).

64. Crystal lattice structures Diamond and graphite

65. T H E E N D Work problems, p. 470-473 #54, 56, 58, 62, 70, 73, 74, 76