1. Bonding: General Concepts Chapter 8

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

3. Bonds Barf Breaking absorbing Release when forming

4. Bond Energy -It is the energy required to break a bond. -It gives us information about the strength of a bonding interaction.

5. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.

6. Pauling Electronegativity Values

7. Three Possible Types of Bonds Nonpolar Covalent (Electrons equally shared.) Polar Covalent (Electrons shared unequally.) Ionic (Electrons are transferred.)

8. 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.

9. Isoelectronic Ions Ions containing the same number of electrons (O 2 , F , Na +, Mg 2+, Al 3+ ) Write the electron configs to all ions above O 2  > F  > Na + > Mg 2+ > Al 3+ largest smallest

10. 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.

11. Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm) When E is positive (+), repulsion is indicated. When E is negative (-), attraction is indicated. This is a statement of Coulomb’s Law where:

12. Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions

13. Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. NaCl, MgCl 2 LiF, LiCl

14. Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released Draw the Lewis Structure for each reactant and product before doing any calculations!

15. Average Bond Dissociation Energies at 298 K BondEnergy, kJ mol -1 C-H414 C-C347 C-Cl377 Cl-Cl243 H-Cl431 The tables above contain information for determining thermodynamic properties of the reaction below. C 2 H 5 Cl(g) + Cl 2 (g)  C 2 H 4 Cl 2 (g) + HCl(g) Calculate the  H  for the reaction above, using the table of average bond dissociation

16. (a)deltaH = energy of bonds broken - energy of bonds formed C 2 H 5 Cl + Cl 2 C 2 H 4 Cl 2 + HCl CH + Cl-Cl - C-Cl + HCl (representing the changes) DH = (414) + 243) - (377 + 431) = -151 kJ

17. Practice +ditto

18. 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 / 2 X 2 (g)  X(g) [endothermic]

20. Formation of an Ionic Solid (continued) 4.Formation of X  ions in the gas phase: X(g) + e   X  (g) [exothermic] 5.Formation of the solid MX lattice energy M + (g) + X  (g)  MX(s) [quite exothermic]

21. Use the following data to estimate ΔHºf for sodium chloride. Na(s) + 1/2Cl2(g) NaCl(s) Lattice energy-786 kJ/mol Ionization energy for Na495 kJ/mol Electron affinity of Cl-349 kJ/mol Bond energy of Cl2239 kJ/mol Enthalpy of sublimation for Na109 kJ/mol

22. Types of Covalent Bonds Polar covalent bond -- covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. A dipole moment exists. Nonpolar covalent bond -- covalent bond in which the electrons are shared equally between both atoms. No dipole moment exists.

23. - covalent bonds are formed by sharing electrons between nuclei. H. +. H ----> H-H - coordinate covalent bonds are bonds where both shared electrons originate on the same atom : NH 3 + H + ----> H + - NH 3

24. Single, Double, & Triple Bonds Single bonds -- one shared pair of electrons. Double bonds -- two shared pairs of electrons. Triple bonds -- three shared pairs of electrons. See bond energy Tables 8.4 & 8.5 on pages 373-374 in Zumdahl.

25. 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.

26. The Effect of an electric field on hydrogen fluoride molecules.

27. Dipole Moment for the water molecule.

28. Dipole moment for the ammonia molecule.

29. Nonpolar molecule--zero dipole moment.

30. 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.

31. Localized Electron 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.

32. 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.

33. Lewis Structures Covalent Compounds Ionic Compounds In ionic compounds, electrons are transferred and ions are formed. In covalent compounds, electrons are shared to form a molecule.

34. Covalent Chalk When writing Lewis structures use (happy-have) / 2 = bonds/distribute rest Cl 2 NH 3 CO 2 CO 3 -2 CO N 2 Sulfur dioxide Nitrite

35. Exceptions When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals. (Central atom) SI4 Br3- Xenon tetrafluoride TeCl4 Xenon Dichloride -

36. Electron Deficient Molecules Beryllium chloride -- BeCl 2 -- is electron deficient with four electrons. It forms a linear molecule. Boron trifluoride -- BF 3 -- is electron deficient with six electrons. It forms a trigonal planar molecule. See page 381 for the reaction between boron trifluoride and ammonia.

37. 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.

38. Exceptions for octet rules ICl3,IF2-XeF4

39. 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 called a resonance hybrid. Ex. the resonance structures for the nitrate ion

40. Draw the Lewis Structure for the following CO 2

41. Finding Formal Charge 1.Valence electrons of the atom (off the table) 2.Electrons owned = # of valence electrons around the atom + the number of bonds. (one electron per bond ) Formal Charge = # valence electrons on the neutral atom – electrons owned by the atom in the resonance structure.

42. SO4 -2 XeO 3

43. Stereochemistry The study of the three- dimensional arrangement of atoms or groups of atoms within molecules and the properties which follow such arrangement.

44. VSEPR Model Valence Shell Electron Pair Repulsion -- The structure around a given atom is determined principally by minimizing electron pair repulsions.

45. Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Put pairs as far apart as possible. 3.Determine positions of atoms from the way electron pairs are shared.(Parent Geometry) 4.Determine the name of molecular structure from positions of the atoms.(Actual Geometry)

46. Molecular Geometry Parent Geometry is electron pair arrangement about the central atom. linear trigonal planar tetrahedral trigonal bipyramidal octahedral Actual Geometry is the arrangement of atoms about the central atom. linear bent trigonal pyramid seesaw T-shaped square pyramid square planar

47. http://www.wwnorton.com/COLLEGE/chemi stry/gilbert/tutorials/chapter_07/vsepr/inde x.html

48. Lone pair of electrons on the ammonia molecule.

49. 08_143 (a)(b) H (c) Lone pair Bonding pair Bonding pair Lone pair H H O O H O Lone pairs on the water molecule.

50. Octahedral structure for phosphorus hexachloride.

51. Octahedral structure for xenon.

52. Parent and actual geometry for xenon tetrafluoride.

53. Three possible arrangements of the electron pairs in triiodide ion.

55. VSEPR Model Summary Determine the Lewis structure(s) for the molecule. For molecules with resonance structures, use any of the structures to predict the molecular structure. Sum the electron pairs around the central atom to determine the parent geometry. The arrangement of the pairs is determined by minimizing electron-pair repulsions.(Actual Geometry)

56. VSEPR Model Summary (Continued) Lone pairs require more space than bonding pairs since they are tightly attracted to only one nucleus. Lone pairs produce slight distortions of bond angles less than 120 o.