1. Chapter 8 Chemical Bonding

2. © 2009, Prentice-Hall, Inc. Chemical Bonds Three basic types of bonds – Ionic Electrostatic attraction between ions – Covalent Sharing of electrons – Metallic Metal atoms bonded to several other atoms due to mobile electrons

3. © 2009, Prentice-Hall, Inc. Energetics of Ionic Bonding It takes 495 kJ/mol to remove electrons from sodium. We get 349 kJ/mol back by giving electrons to chlorine.

4. © 2009, Prentice-Hall, Inc. Energetics of Ionic Bonding But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic! What is as yet unaccounted for is the electrostatic attraction between the newly- formed sodium cation and chloride anion.

5. © 2009, Prentice-Hall, Inc. Energetics of Ionic Bonding There must be a third piece to the puzzle. What is as yet unaccounted for is the electrostatic attraction between the newly- formed sodium cation and chloride anion.

6. © 2009, Prentice-Hall, Inc. Lattice Energy This third piece of the puzzle is the lattice energy: – The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. The energy associated with electrostatic interactions is governed by Coulomb’s law: E el =  Q1Q2dQ1Q2d Lattice energy increases with the charge on the ions. It also increases with decreasing size of ions. (decreased distance between the ions)

7. © 2009, Prentice-Hall, Inc. Explain the difference in the lattice energies of KF, CaO and ScN Explain the difference in the lattice energies of LiF, NaCl and KBr

8. © 2009, Prentice-Hall, Inc. Covalent Bonding In covalent bonds atoms share electrons. There are several electrostatic interactions in these bonds: – Attractions between electrons and nuclei – Repulsions between electrons – Repulsions between nuclei

9. © 2009, Prentice-Hall, Inc. Polar Covalent Bonds Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine has a greater attraction for the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

10. © 2009, Prentice-Hall, Inc. Electronegativity Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. On the periodic chart, electronegativity increases as you go… – …from left to right across a row. – …from the bottom to the top of a column.

11. © 2009, Prentice-Hall, Inc. Polar Covalent Bonds When two atoms share electrons unequally, a bond dipole results. The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:  = Q r It is measured in debyes (D).

12. © 2009, Prentice-Hall, Inc. Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.

13. © 2009, Prentice-Hall, Inc. Writing Lewis Structures 1.Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. – If it is an anion, add one electron for each negative charge. – If it is a cation, subtract one electron for each positive charge. PCl 3 5 + 3(7) = 26

14. © 2009, Prentice-Hall, Inc. Writing Lewis Structures 2.The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. 3.Fill the octets of the outer atoms. Keep track of the electrons: 26 - 6 = 20

15. © 2009, Prentice-Hall, Inc. Writing Lewis Structures 4.Fill the octet of the central atom. Keep track of the electrons: 26 - 6 = 20; 20 - 18 = 2; 2 - 2 = 0

16. © 2009, Prentice-Hall, Inc. Writing Lewis Structures 5.If you run out of electrons before the central atom has an octet… …move lone pairs to form multiple bonds until it does.

17. Writing Lewis Structures 6. If too many electrons remain, assign them to the central atom, even if it would exceed an octet.

18. © 2009, Prentice-Hall, Inc. Writing Lewis Structures Then assign formal charges. – For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. – Subtract that from the number of valence electrons for that atom: the difference is its formal charge. F.C. = valence e - - (dots+ dashes)

19. © 2009, Prentice-Hall, Inc. Writing Lewis Structures The best Lewis structure… – …is the one with the fewest charges. – …puts a negative charge on the most electronegative atom.

20. Use formal charge to decide the best structure to represent the thiocyanate ion: and for Nitrogen (II) oxide:

21. © 2009, Prentice-Hall, Inc. Resonance This is the Lewis structure we would draw for ozone, O 3. - + Until now, we have not been concerned with the location of the double bond. We could also draw this structure:

22. © 2009, Prentice-Hall, Inc. Resonance But this is at odds with the true, observed structure of ozone, in which… – …both O-O bonds are the same length and strength – …both outer oxygens have a formal charge of -1/2.

23. © 2009, Prentice-Hall, Inc. Resonance One Lewis structure cannot accurately depict a molecule like ozone. We use multiple structures, resonance structures, to describe the molecule. In fact, the electrons that form the second O-O bond in the double bond are not found in just one location… they are delocalized.

24. © 2009, Prentice-Hall, Inc. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

25. © 2009, Prentice-Hall, Inc. Resonance The organic compound benzene, C 6 H 6, has two resonance structures. It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

26. © 2009, Prentice-Hall, Inc. Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: – Ions or molecules with an odd number of electrons – Ions or molecules with less than an octet – Ions or molecules with more than eight valence electrons (an expanded octet)

27. © 2009, Prentice-Hall, Inc. Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. Nitrogen can form such molecules… they are sometimes referred to as free radicals and may negatively impact human health.

28. © 2009, Prentice-Hall, Inc. Fewer Than Eight Electrons Consider BF 3 : – Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. – This would not be an accurate picture of the distribution of electrons in BF 3.

29. © 2009, Prentice-Hall, Inc. Fewer Than Eight Electrons Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

30. © 2009, Prentice-Hall, Inc. More Than Eight Electrons The only way PCl 5 can exist is if phosphorus has 10 electrons around it. It is allowed to expand its octet. – It is assumed that d orbitals can participate in bonding. Atoms that have d orbitals are found in periods 3 through 7… and can have an expanded octet.

31. © 2009, Prentice-Hall, Inc. Covalent Bond Strength Most simply, the strength of a bond is measured by determining how much energy is required to break the bond. This is the bond enthalpy. The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol.

32. © 2009, Prentice-Hall, Inc. Average Bond Enthalpies This table lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an endothermic process.

33. © 2009, Prentice-Hall, Inc. Enthalpies of Reaction Yet another way to estimate  H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed. In other words,  H rxn =  (bond enthalpies of bonds broken) -  (bond enthalpies of bonds formed)

34. Enthalpies of Reaction © 2009, Prentice-Hall, Inc. CH 4 (g) + Cl 2 (g)  CH 3 Cl (g) + HCl (g) In this example, one C-H bond and one Cl-Cl bond are broken; one C-Cl and one H-Cl bond are formed. Calculate the energy absorbed in breaking bonds. Calculate the energy released when bonds form. Find the difference between them and assign the appropriate sign (- or +)

35. © 2009, Prentice-Hall, Inc. Bond Enthalpy and Bond Length We can also measure an average bond length for different bond types. As the number of bonds between two atoms increases, the bond length decreases.

36. Chapter 9 Molecular Geometries and Bonding Theories

37. © 2009, Prentice-Hall, Inc. What Determines the Shape of a Molecule? Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

38. © 2009, Prentice-Hall, Inc. Electron Domains We can refer to the electron pairs as electron domains. In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. The central atom in this molecule, A, has four electron domains.

39. © 2009, Prentice-Hall, Inc. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

40. © 2009, Prentice-Hall, Inc. Electron-Domain Geometries These are the electron-domain geometries for two through six electron domains around a central atom.

41. © 2009, Prentice-Hall, Inc. Electron-Domain Geometries All one must do is count the number of electron domains in the Lewis structure. The geometry will be that which corresponds to the number of electron domains.

42. © 2009, Prentice-Hall, Inc. Molecular Geometries The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

43. © 2009, Prentice-Hall, Inc. Linear Electron Domain In the linear domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

44. © 2009, Prentice-Hall, Inc. Trigonal Planar Electron Domain There are two molecular geometries: – Trigonal planar, if all the electron domains are bonding, – Bent, if one of the domains is a nonbonding pair.

45. © 2009, Prentice-Hall, Inc. Tetrahedral Electron Domain There are three molecular geometries: – Tetrahedral, if all are bonding pairs, – Trigonal pyramidal if one is a nonbonding pair, – Bent if there are two nonbonding pairs.

46. © 2009, Prentice-Hall, Inc. Trigonal Bipyramidal Electron Domain There are two distinct positions in this geometry: – Axial – Equatorial

47. © 2009, Prentice-Hall, Inc. Trigonal Bipyramidal Electron Domain Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

48. © 2009, Prentice-Hall, Inc. Octahedral Electron Domain All positions are equivalent in the octahedral domain. There are three molecular geometries: – Octahedral – Square pyramidal – Square planar

49. © 2009, Prentice-Hall, Inc. Larger Molecules In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole. This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

50. © 2009, Prentice-Hall, Inc. Polarity Molecules whose centers of negative and positive charge do not coincide are said to be polar molecules. Non-polar molecules have a zero dipole moment.

51. © 2009, Prentice-Hall, Inc. Overlap and Bonding We think of covalent bonds forming through the sharing of electrons by adjacent atoms. In such an approach this can only occur when orbitals on the two atoms overlap.

52. © 2009, Prentice-Hall, Inc. Hybrid Orbitals But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.

53. © 2009, Prentice-Hall, Inc. Hybrid Orbitals Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals. – These sp hybrid orbitals have two lobes like a p orbital. – One of the lobes is larger and more rounded as is the s orbital. These two degenerate orbitals would align themselves 180  from each other…. a linear geometry!

54. © 2009, Prentice-Hall, Inc. Hybrid Orbitals When an s orbital and 2 p orbitals are mixed, we get…… …three degenerate sp 2 orbitals.

55. © 2009, Prentice-Hall, Inc. Hybrid Orbitals Once you know the electron-domain geometry, you know the hybridization state of the atom. The number of electron domains should equal the number of hybrid orbitals.

56. © 2009, Prentice-Hall, Inc. Sigma (  ) Bonds Sigma bonds are characterized by – Head-to-head overlap. – Cylindrical symmetry of electron density about the internuclear axis.

57. © 2009, Prentice-Hall, Inc. Pi (  ) Bonds Pi bonds are characterized by – Side-to-side overlap. – Electron density above and below the internuclear axis.

58. © 2009, Prentice-Hall, Inc. Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering. In a multiple bond one of the bonds is a  bond and the rest are  bonds.

59. © 2009, Prentice-Hall, Inc. Delocalized Electrons: Resonance When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

60. © 2009, Prentice-Hall, Inc. Delocalized Electrons: Resonance In reality, each of the four atoms in the nitrate ion has a p orbital. The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen. This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.