Bonding

1. Bonding 8 out of 75 M/C Questions Free Response—Every year

2. In General: • All bonds occur because of electrostatic attractions. • Formation of molecules and the state of matter of a substance depends on the attractions between electron clouds of one atom and nucleus of another atom.

3. Bonding—General Rules • A metal and a nonmetal bond so that charges on the ions cancel. • When two nonmetals react to form a covalent bond, they share electrons in a way that gives both atoms a Noble gas configuration.

4. Bonding—General Rules • nonmetal + representative metal binary ionic compound: ions form to give the nonmetal the valence electron configuration of the next noble gas atom, and valence orbitals of the metal are emptied

5. Bonding—General Rules • Most bonds are combinations of ionic/covalent character • The more different the atoms bonding are, the more ionic character of the bond. • More similar—more covalent

6. Covalent Sharing Molecules Structural formula Ionic Transfer / charged Compounds Formula Unit Word Association

7. Coulomb’s Law • Describes energy of interaction between ions • E = k J*nm(Q1Q2 / r) • Q1 & Q2numerical ion charges • r distance between ion centers • Negative ans. means ion pair is more stable than individual ions.

8. Coulomb’s Law Simplified • As charges on ions increase, the energy of the bond increases. • As the radius of the atoms increases, the energy of the bond decreases.

9. Which bond has greater energy? • NaF or MgO? • NaCl or CsI? • MgCl2 or NaCl? • MgCl2 or BaBr2?

10. Bond Length • Distance between bonding atoms at which energy is minimized • Atoms position themselves to minimize repulsions and maximize attraction & thus achieve lowest possible energy.

11. Predicting Formulas for Ionic Compounds • Metal—positive charge equal to # of valence electrons • Nonmetal—negative charge equal to # of electrons away from next Noble gas • Compound—charges must cancel

12. Predict Formulas for: • Potassium sulfide • Barium chloride • Aluminum oxide • Magnesium phosphide

13. Exceptions to the Rule: • Sn—forms both +2 and +4 ions • Pb—forms both +2 and +4 ions • Bi—forms +3 and +5 ions • Tl—forms +1 and +3 ions  “no simple explanation for this behavior”

14. Lattice Energy • Change in energy taking place when separated gaseous ions are packed together to form an ionic solid • Energy released when an ionic solid forms from its ions

15. Calculating Energy Change • The sum of energy changes in each step • Formation of lithium fluoride: • Sublimation of solid lithium: 161 kJ per mole • Ionization of lithium: 520 kJ per mole

16. Calculating Energy Change • Dissociation of F2: 77 kJ per mole of F atoms • Ionization of F: -328 kJ per mole (Electron affinity of F) • Formation of LiF: -1047 kJ per mole

17. Calculating Energy Change • Sum: 161 + 520 + 77 – 328 – 1047 = -617 kJ per mole LiF (negative sign means process is exothermic & product has lower energy)

18. Calculating Lattice Energy • Lattice energy can be calculated using a form of Coulomb’s Law: Lattice energy = k(Q1Q2 / r) k = proportionality constant dependent on solid’s structure & electron configuration of the solid

19. P. 365—energy changes in forming LiF and MgO

20. Lattice Energy • Much larger changes in Mg and O because of greater Q1 and Q2 • Even though Mg and O must ionize twice (which requires lots of energy) the overall change is very exothermic

21. DH for Covalent Compunds • Breaking bonds requires energy (+). • Forming bonds releases energy (-). • Determine which bonds break and which bonds form. • Add values for broken bonds. • Subtract values for formed bonds.

22. Find DH CH4 + Cl2 + F2 CF2Cl2 + 2HF + 2HCl

23. Polar Covalent Bonds • Most bonds have both ionic and covalent character • Polar covalent bonds—unequally shared electrons • Not different enough to be purely ionic nor similar enough to be purely covalent

24. Polar Covalent Bonds • Strength of polarity depends on electronegativity (ability of an atom to attract electrons to itself). • Higher electronegativity results in higher negative charge. • Partial charge results in a stronger bond & stronger inter-molecular attraction.

25. Dipoles • + indicates direction of dipole—points to negative end • Dipolar molecules orient themselves one way in an electric field. • Dipoles can cancel.

26. Canceling Dipoles • If evenly spacedidentical bonds exist, dipoles cancel each other. • Linear—2 identical bonds separated by 180* (CO2) • Trigonal planar—3 identical bonds separated by 120* (SO3) • Tetrahedral—4 identical bonds separated by 109.5* (CH4)

27. Ionic Bonds with Covalent Character • No totally ionic bonds exist. • Percent ionic character can be calculated: (measured dipole moment of X-Y) (calculated dipole moment of X+Y-) X 100%

28. Ionic Bonds with Covalent Character • X-Y represents the molecule • X+Y- represents the ions • Graph—p. 367—Bonds with more than 50% ionic character are classified as ionic

29. New Definition • Ionic compound—any substance that conducts electricity in its liquid state (melted, not dissolved)

30. Beyond the Bond—Intermolecular Forces & States of Matter

31. Intermolecular Forces • Weaker than chemical bond • Affect structure and state of matter

32. Dipole-Dipole Forces • Positive and negative ends of polar molecules attract each other. • About 1% as strong as covalent or ionic bonds • Weaken as distance between molecules increases

33. Hydrogen Bonding • Especially strong dipole-dipole force • Occurs when H bonds to a strongly electronegative atom—O, N, or F • Very strong because 1) molecule is very polar & 2) small size of H

34. H Bonding • Example—water • More pronounced in molecules formed from small atoms (dipoles can come closer) • High boiling point

35. London Dispersion Forces • Forces that exist in all substances but are important only in Noble gases and nonpolar molecules • Result from temporary dipoles formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atom • VERY WEAK

36. Bonding models

37. Problems with Current Model • If s & p orbitals are different, bonds formed from them should be different. • Since p orbitals are perpendicular to one another, we would expect 90o bond angles. • Neither of these things is true.

38. Hybridization • Explains discrepancies • The mixing of native atomic orbitals to form special orbitals for bonding • Hybrid orbitals have shapes and energies that are between those of the native orbitals.

39. Example: CH4 • Four hydrogen atoms combine with carbon in each of four orbitals. The one s and three p orbitals hybridize to form four hybrid sp3 orbitals.

40. Hybridization and Geometry • The number of areas of charge density (bonded atoms + unshared pairs) relates to both hybridization and geometry. • sp3—4 areas/orbitals—tetrahedral • sp2—3 areas/orbitals—trigonal planar • sp—2 areas/orbitals—linear

41. sp2 A molecule such as BH3 or an ion such as NO3- exhibit sp2 hybridization. This results in a trigonal planar geometry.

42. Multiple Bonds • First bond in a hybrid orbital (sigma) s • Second or third in an unhybridized orbital (pi) p

43. Bonding in Ethene

44. sp • One s and one p orbital combine to form sp hybrids that will be at 90o to one another as in BeF2or CO2.

45. dsp3 • Five electron pairs require the use of a d orbital. Only elements in level 3 or above can do this as in PF5 which has trigonal bipyramidal geometry.

46. d2sp3 • The six electron pairs in a compound such as SF6 result in the use of two d orbitals and octahedral geometry.

47. Localized Electron Model • Assumes electrons are located between two atoms • Includes several parts: • Description of valence electron arrangement (Lewis Structure) • Prediction of Geometry (VSEPR) • Description of types of orbitals (Hybridization)

48. Shortcomings • Localized electron model • Assumes electrons are localized so resonance must be added. • Does not adequately address unpaired electrons. • Gives no information about bond energies.

49. Molecular Orbital Model • Uses molecular orbitals (instead of atomic orbitals) to address shortcomings of localized electron model.

50. Molecular Orbitals • Combining atoms forms two molecular orbitals—not one atomic orbital.