1. Chemical Bonding and Molecular Architecture Structure and Shapes of Chemicals

2. Bonds Forces that hold groups of atoms together and make them function as a unit.

3. Bond Energy It is the energy required to break or released in making a bond. It gives us information about the strength of a bonding interaction. Ionic bonds—strong attractions between oppositely charged ions Covalent bonds—attraction between non-metal atoms as both atoms share electrons

4. Bond Length The distance where the system energy is a minimum.

6. Ionic Bonds -Formed from electrostatic attractions of closely packed, oppositely charged ions. -Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

7. Ionic Configuration and Size Ions are formed when electrons are gained or lost from an atom. The gain or loss follows the pattern called the “octet rule”, that an atom forms an ion in which it attains the same electron configuration as the nearest noble gas. Most metals therefore lose electrons, and as a result get smaller. The trend is “greater +, smaller size.” Likewise, nonmetals gain electrons to form ions, thus increasing in size by the opposite rule to metals.

9. Isoelectronic Ions Ions containing the the same number of electrons, due to attaining the configuration of the same noble gas (O 2 , F , Na +, Mg 2+, Al 3+ ) All attain to Ne O 2  > F  > Na + > Mg 2+ > Al 3+ largest smallest

10. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Periodic trend –increases across the table to the halogen column. Decreases down a group. Least at Cs (0.7), greatest at F (4.0).

12. Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. Polar bonds shown as arrow with point toward negative pole, + toward the positive pole

13. Electronegativity and Polarity of Bonds Subtract lower EN from higher EN Difference % Ionic Character Type of Bond 00Nonpolar Covalent 0.1-0.51-5%Slightly polar covalent 0.1-0.51-5%Slightly polar covalent 0.6-1.56-40%Polar Covalent 0.6-1.56-40%Polar Covalent > 1.5over 40% Ionic > 1.5over 40% Ionic Compounds with over 50% ionic character are considered to be totally ionic solids. These compounds are often called salts.

16. Homework!! p. 395ff 11, 14, 15, 20

17. Achieving Noble Gas Electron Configurations (NGEC) Two nonmetals react: They share electrons to achieve NGEC. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

18. Binary Ionic--Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M + (g) + X  (g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

19. Formation of an Ionic Solid 1.Sublimation of the solid metal M(s)  M(g) [endothermic] 2.Ionization of the metal atoms M(g)  M + (g) + e  [endothermic] 3.Dissociation of the nonmetal 1 /2X 2 (g)  X(g) [endothermic] 4.Formation of X  ions in the gas phase: X(g) + e   X  (g) [exothermic] 5.Formation of the solid MX M + (g) + X  (g)  MX(s) [quite exothermic]

21. Covalent Chemical Bonds Happen when collections of atoms are more stable than the separate atoms. They provide a method for dividing up energy when stable molecules are formed from atoms. Covalent bonds are due to shared electron pairs. One pair shared is a single bond, two makes a double bond, three make a triple bond. As bond order increases (single, double, triple), bond length shortens

22. Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released

24. Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Two types of electron pairs: bonding pairs and lone pairs. Bonding pairs are linkages between atoms, lone pairs are electrons solely owned by an atom.

25. Localized Electron Model Elements of the Model 1.Description of valence electron arrangement (Lewis structure). 2.Prediction of geometry (VSEPR model). 3.Description of atomic orbital types used to share electrons or hold lone pairs.

26. Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

27. Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

28. Rules for Drawing Lewis Structures Add up all of the valence electrons for the atoms involved in the molecule Select a most likely central atom and arrange other atoms around it. Place pairs of electrons between atoms. Arrange the remaining electrons around external atoms first. If the central atom is not satisfied, form double or triple bonds to make the molecule work.

30. Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

33. Homework p. 397ff 31, 36, 39, 42, 50, 57

34. Molecular Architecture The structure of a molecule is important in how it reacts and to its physical properties Once the Lewis structure of a molecule is determined, the shape of the molecule then can be predicted according to the VSEPR model.

35. VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

37. Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Count pairs, both bonding and lone pairs around the central atom. 3.Determine positions of atoms from the way electron pairs are shared. 4.Determine the name of molecular structure from the number of bonding and lone pairs and their necessary arrangements. Remember that lone pairs prefer to be at 120º or greater from each other.

41. Homework!! p. 399ff 59, 62, 73, 78, 79, 91

42. Hybridization The mixing of atomic orbitals to form special orbitals for bonding. The atoms are responding as needed to give the minimum energy for the molecule. To determine hybridization, count lone and bonding pairs, but count multiple bonds only once.

48. A sigma (  ) bond centers along the internuclear axis. A pi (  ) bond occupies the space above and below the internuclear axis.

53. The Localized Electron Model -Draw the Lewis structure(s) -Determine the arrangement of electron pairs (VSEPR model). -Specify the necessary hybrid orbitals.

56. Homework p. 432ff 5, 8, 11

57. Molecular Orbitals (MO) Analagous to atomic orbitals for atoms, MOs are the quantum mechanical solutions to the organization of valence electrons in molecules.

58. Types of MOs bonding: lower in energy than the atomic orbitals from which it is composed. antibonding: higher in energy than the atomic orbitals from which it is composed.

61. Bond Order (BO) Difference between the number of bonding electrons and number of antibonding electrons divided by two.

62. Paramagnetism -unpaired electrons -attracted to induced magnetic field -much stronger than diamagnetism

63. Outcomes of MO Model 1.As bond order increases, bond energy increases and bond length decreases. 2.Bond order is not absolutely associated with a particular bond energy. 3.N 2 has a triple bond, and a correspondingly high bond energy. 4.O 2 is paramagnetic. This is predicted by the MO model, not by the LE model, which predicts diamagnetism.

65. Combining LE and MO Models  bonds can be described as being localized.  bonding must be treated as being delocalized.

67. Homework p. 434 ff 17, 22, 25, 37