1. Bonding

2. What is a Bond? l A force that holds atoms together. l Why? l We will look at it in terms of energy. l Bond energy the energy required to break a bond. l Why are compounds formed? l Because it gives the system the lowest energy.

3. Ionic Bonding l An atom with a low ionization energy reacts with an atom with high electron affinity. l The electron moves. l Opposite charges hold the atoms together.

4. Coulomb's Law l E= 2.31 x 10 -19 J · nm(Q 1 Q 2 )/r l Q is the charge. l r is the distance between the centers. l If charges are opposite, E is negative l exothermic l Same charge, positive E, requires energy to bring them together.

5. What about covalent compounds? l The electrons in each atom are attracted to the nucleus of the other. l The electrons repel each other, l The nuclei repel each other. l The reach a distance with the lowest possible energy. l The distance between is the bond length.

6. 0 Energy Internuclear Distance

7. 0 Energy Internuclear Distance

8. 0 Energy Internuclear Distance

9. 0 Energy Internuclear Distance

10. 0 Energy Internuclear Distance Bond Length

11. 0 Energy Internuclear Distance Bond Energy

12. Covalent Bonding l Electrons are shared by atoms. l These are two extremes. l In between are polar covalent bonds. l The electrons are not shared evenly. l One end is slightly positive, the other negative. Indicated using small delta 

13. H - F ++ --

14. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

15. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- + -

16. ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- - +

17. Electronegativity l The ability of an electron to attract shared electrons to itself. l Pauling method l Imaginary molecule HX l Expected H-X energy = H-H energy + X-X energy 2  = (H-X) actual - (H-X) expected

18. Electronegativity  is known for almost every element l Gives us relative electronegativities of all elements. l Tends to increase left to right. l decreases as you go down a group. l Noble gases aren’t discussed. l Difference in electronegativity between atoms tells us how polar.

19. Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases

20. Dipole Moments l A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), l or has a dipole moment. l Center of charge doesn’t have to be on an atom. l Will line up in the presence of an electric field.

21. How It is drawn H - F ++ --

22. Which Molecules Have Them? l Any two atom molecule with a polar bond. l With three or more atoms there are two considerations. l There must be a polar bond. l Geometry can’t cancel it out.

23. Geometry and polarity l Three shapes will cancel them out. l Linear

24. Geometry and polarity l Three shapes will cancel them out. l Planar triangles 120º

25. Geometry and polarity l Three shapes will cancel them out. l Tetrahedral

26. Geometry and polarity l Others don’t cancel l Bent

27. Geometry and polarity l Others don’t cancel l Trigonal Pyramidal

28. Ions l Atoms tend to react to form noble gas configuration. l Metals lose electrons to form cations l Nonmetals can share electrons in covalent bonds. l When two non metals react.(more later) l Or they can gain electrons to form anions.

29. Ionic Compounds l We mean the solid crystal. l Ions align themselves to maximize attractions between opposite charges, l and to minimize repulsion between like ions. l Can stabilize ions that would be unstable as a gas. l React to achieve noble gas configuration

30. Size of ions l Ion size increases down a group. l Cations are smaller than the atoms they came from. l Anions are larger. l across a row they get smaller, and then suddenly larger. l First half are cations. l Second half are anions.

31. Periodic Trends l Across the period nuclear charge increases so they get smaller. l Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

32. Size of Isoelectronic ions l Iso - same l Iso electronic ions have the same # of electrons l Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 l All have 10 electrons. l All have the configuration 1s 2 2s 2 2p 6

33. Size of Isoelectronic ions l Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

34. Forming Ionic Compounds l Lattice energy - the energy associated with making a solid ionic compound from its gaseous ions. M + (g) + X - (g)  MX(s) l This is the energy that “pays” for making ionic compounds. l Energy is a state function so we can get from reactants to products in a round about way.

35. Na(s) + ½F 2 (g)  NaF(s) First sublime NaNa(s)  Na(g)  H = 109 kJ/mol Ionize Na(g) Na(g)  Na + (g) + e -  H = 495 kJ/mol Break F-F Bond½F 2 (g)  F(g)  H = 77 kJ/mol Add electron to F F(g) + e -  F - (g)  H = -328 kJ/mol

36. Na(s) + ½F 2 (g)  NaF(s) Lattice energy Na(s) + ½F 2 (g)  NaF(s)  H = -928 kJ/mol

37. Calculating Lattice Energy l Lattice Energy = k(Q 1 Q 2 / r) l k is a constant that depends on the structure of the crystal. l Q’s are charges. l r is internuclear distance. l Lattice energy is greater with more highly charged ions.

38. Calculating Lattice Energy l This bigger lattice energy “pays” for the extra ionization energy. l Also “pays” for unfavorable electron affinity.

39. Partial Ionic Character l There are probably no totally ionic bonds between individual atoms. l Calculate % ionic character. l Compare measured dipole of X-Y bonds to the calculated dipole of X + Y - the completely ionic case. l % dipole = Measured X-Y x 100 Calculated X + Y - l In the gas phase.

40. % Ionic Character Electronegativity difference 25% 50% 75%

41. How do we deal with it? l If bonds can’t be ionic, what are ionic compounds? l And what about polyatomic ions? l An ionic compound will be defined as any substance that conducts electricity when melted. l Also use the generic term salt.

42. The Covalent Bond l The forces that cause a group of atoms to behave as a unit. l Why? l Due to the tendency of atoms to achieve the lowest energy state. l It takes 1652 kJ to dissociate a mole of CH 4 into its ions l Since each hydrogen is hooked to the carbon, we get the average energy = 413 kJ/mol

43. l CH 3 Cl has 3 C-H, and 1 C - Cl l the C-Cl bond is 339 kJ/mol l The bond is a human invention. l It is a method of explaining the energy change associated with forming molecules. l Bonds don’t exist in nature, but are useful. l We have a model of a bond.

44. What is a Model? l Explains how nature operates. l Derived from observations. l It simplifies the and categorizes the information. l A model must be sensible, but it has limitations.

45. Properties of a Model l A human inventions, not a blown up picture of nature. l Models can be wrong, because they are based on speculations and oversimplification. l Become more complicated with age. l You must understand the assumptions in the model, and look for weaknesses. l We learn more when the model is wrong than when it is right.

46. Covalent Bond Energies l We made some simplifications in describing the bond energy of CH 4 l Each C-H bond has a different energy. CH 4  CH 3 + H  H = 435 kJ/mol CH 3  CH 2 + H  H = 453 kJ/mol CH 2  CH + H  H = 425 kJ/mol CH  C + H  H = 339 kJ/mol l Each bond is sensitive to its environment.

47. Averages l Have made a table of the averages of different types of bonds pg. 365 l single bond one pair of electrons is shared. l double bond two pair of electrons are shared. l triple bond three pair of electrons are shared. l More bonds, shorter bond length.

48. Using Bond Energies We can find  H for a reaction. l It takes energy to break bonds, and end up with atoms (+). l We get energy when we use atoms to form bonds (-). If we add up the energy it took to break the bonds, and subtract the energy we get from forming the bonds we get the  H. l Energy and Enthalpy are state functions.

49. Find the energy for this 2 CH 2 = CHCH 3 + 2NH 3 O2O2 +  2 CH 2 = CHC  N + 6 H 2 O C-H 413 kJ/mol C=C 614kJ/mol N-H 391 kJ/mol O-H 467 kJ/mol O=O 495 kJ/mol C  N 891 kJ/mol C-C 347 kJ/mol

50. Localized Electron Model l Simple model, easily applied. l A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. l Three Parts 1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals (Chapt 9)

51. Lewis Structure l Shows how the valence electrons are arranged. l One dot for each valence electron. l A stable compound has all its atoms with a noble gas configuration. l Hydrogen follows the duet rule. l The rest follow the octet rule. l Bonding pair is the one between the symbols.

52. Rules l Sum the valence electrons. l Use a pair to form a bond between each pair of atoms. l Arrange the rest to fulfill the octet rule (except for H and the duet). lH2OlH2O l A line can be used instead of a pair.

53. A useful equations l (happy-have) / 2 = bonds l POCl 3 P is central atom l SO 4 -2 S is central atom l SO 3 -2 S is central atom l PO 4 -2 S is central atom l SCl 2 S is central atom

54. Exceptions to the octet l BH 3 l Be and B often do not achieve octet l Have less than and octet, for electron deficient molecules. l SF 6 l Third row and larger elements can exceed the octet l Use 3d orbitals? l I 3 -

55. Exceptions to the octet l When we must exceed the octet, extra electrons go on central atom. l ClF 3 l XeO 3 l ICl 4 - l BeCl 2

56. Resonance l Sometimes there is more than one valid structure for an molecule or ion. l NO 3 - l Use double arrows to indicate it is the “average” of the structures. l It doesn’t switch between them. l NO 2 - l Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

57. Formal Charge l For molecules and polyatomic ions that exceed the octet there are several different structures. l Use charges on atoms to help decide which. l Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

58. Formal Charge l The difference between the number of valence electrons on the free atom and that assigned in the molecule. l We count half the electrons in each bond as “belonging” to the atom. l SO 4 -2 l Molecules try to achieve as low a formal charge as possible. l Negative formal charges should be on electronegative elements.

59. Examples l XeO 3 l NO 4 -3 l SO 2 Cl 2

60. VSEPR l Lewis structures tell us how the atoms are connected to each other. l They don’t tell us anything about shape. l The shape of a molecule can greatly affect its properties. l Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

61. VSEPR l Molecules take a shape that puts electron pairs as far away from each other as possible. l Have to draw the Lewis structure to determine electron pairs. l bonding l nonbonding lone pair l Lone pair take more space. l Multiple bonds count as one pair.

62. VSEPR l The number of pairs determines –bond angles –underlying structure l The number of atoms determines –actual shape

63. VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal

64. Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

65. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

66. Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

67. No central atom l Can predict the geometry of each angle. l build it piece by piece.

68. How well does it work? l Does an outstanding job for such a simple model. l Predictions are almost always accurate. l Like all simple models, it has exceptions.