Ch 9 Bonding and Molecular Structure: Fundamental Concepts

1. Ch 9 Bonding and Molecular Structure: Fundamental Concepts • Valence Electrons • Chemical Bond Formation • Bonding in Ionic Compounds • Covalent Bonding and Lewis Structures • Resonance • Exceptions to the Octet Rule • Charge Distribution in Covalent Molecules • Bond Properties • Molecular Shapes • Molecular Polarity

2. Valence Electrons • There are core electrons and valence electrons • Valence electrons, outermost shell,determine chemical properties • Main group elements - valance electrons • are the s and p, • Same as the group number • Transition elements - valance electrons are the ns and the (n-1) d orbitals.

3. Valence Electrons • Lewis electron dot diagrams • The element’s symbol represents the atomic nucleus and the core electrons. • Up to 4 valence electrons, dots, are placed around the symbol • Any remaining electrons are paired with those 4 • A full set of 8 electrons is an octet and stable

4. Lewis electron dot diagrams Covalent bonding involves sharing of valence electrons Between atoms.

5. Bonding in Ionic Compounds Formation of these compounds is exothermic in part ionization Of Na and Ca is low.

6. Bonding in Ionic Compounds

7. Bonding in Ionic Compounds

8. Bonding in Ionic Compounds

9. Lattice Energy

10. Lattice Energy Li+ 1036 853 807 757 Na+ 923 787 747 704 K+ 821 715 682 649 Rb+ 785 689 660 630 Cs+ 740 659 631 604 The bond between ions of opposite charge is strongest when the ions are small.

11. Lattice Energy The ionic bond should also become stronger as the charge on the ions becomes larger. The data in the table below show that the lattice energies for salts of the OH- and O2- ions increase rapidly as the charge on the ion becomes larger. OH- O2- Na+900 2481 Mg2+ 3006 3791 Al 3+ 5627 15,916

12. Lattice Energy

13. Lattice Energy • Why do NaCl2 or Na2Cl not exist? • The energy for Na+ to lose a 2nd electron is very high and endothermic. • So the energy of the reaction would not be negative • The reaction would not be product favored • Cl- is not favorable to add a second electron • The amount of energy required is very large

14. Covalent Bonding and Lewis Structure • Molecules or polyatomic ions made of nonmetal atoms bonding covalently. Bonding Pair Shared Pair each 2 shared dots is one stick bond Nonbonding Pair, Unshared Pair, Lone Pair Double Bonds 2 shared pair 4 shared electrons 2 stick bonds Triple Bonds 3 shared pair 6 shared electrons 3 stick bonds

15. Drawing Lewis Structures http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/lewis.html misterguch.brinkster.net/lewisstructures.html http://pages.towson.edu/ladon/lewis.html

16. Drawing Lewis Structures • Decide on the Central Atom the one with the lowest electron affinity (or most bonds) • Determine the total number of valence electrons in the molecule or ion -add up all electron dots and charges • Draw skeletal structure with single bonds • Use remaining electrons to surround terminal atoms with 8 dots (except H) • If central atom has fewer than 8 electrons, move lone pair to create multiple bonds and allow all atoms to have 8 electrons.

17. Drawing Lewis Structures • Practice drawing structures for NH4+ CO NO+ SO4 2-

18. Drawing Lewis Structures • Practice drawing structures for CH4C 2 H6 C 2 H4 C 2 H2 NH2- H3O+

19. Lewis Structures of Acids and their Anions

20. Lewis Structures of Acids and their Anions

21. Isoelectronic Species

22. Isoelectronic Species

23. Resonance Structure When two equivalent structures can be drawn for one molecule, the extra bond was believed to resonate (flip) between the two. If that were occurring,t he lengths of the bonds would be different on each side. This difference in bond length is not observed. Resonance theory was proposed by Linus Pauling as a composite. The resonance Hybrid will provide two equal bonds with bond lengths

24. Resonance Structure

25. Resonance Structure

26. Exceptions to the Octet Rule • Atoms that have fewer than four pair of electrons on a central atom • Generally Group 2A or 3A • Classic example BF3 • Non octet structures • Atoms that have more than four pair of electrons on a central atom. • Generally Group 4A, 5A, 6A, 7A, or 8 • (Yes, Group 8 can form bonds) • Expanded Octet

27. Expanded Octet Only elements of the third or higher periods in the periodic Table form compounds and ions in which an octet is exceeded. Second period elements have only s and p orbitals Elements win the third and higher periods, the d orbitals in the Outer shell are traditionally included among valence orbitals For the elements. The extra orbitals provide the elements with an opportunity to Accommodate up to 12 electrons.

28. Expanded Octet

29. Molecules with an Odd Number of Electrons • There is a small group of molecules with an odd number of electrons • Molecules with an odd number of electrons do not obey the Octet • The single electron classifies the molecules as a free radical. • Free radicals are more reactive than molecules with paired electrons

30. Charge Distribution in the Bonds &Molecules • Lewis structures indicate that electrons are evenly distributed along bonds. • Valence electrons are not distributed among the atoms as evenly as suggested. • Some atoms may have a slight negative charge or a slight positive charge • Electrons are drawn more strongly to one atom in the bond. • Electron distribution defines the uneven distribution of electrons.

31. Charge Distribution in the Covalent Bonds &Molecules • Electrons in covalent bonds are not necessarily evenly distributed. • Charge distribution defines the uneven distribution of electrons. • Polar diatomic molecules will line up by polarity • Determines the site at which reactions occur (like the attachment of an H+), will the H attach to the O or Cl end of a OCl-

32. Formal Charges on Atoms

33. Formal Charges on Atoms

34. Formal Charges on Atoms Calculate formal charges for the atoms in NH4+ CN - CO32- SO3

35. Bond Polarity and Electronegativity • Pure covalent equal sharing-few molecules • Polar covalent bonds - electron pair shared unequally molecule acts as a dipole • 1930 Linus Pauling determined Electronegativity • The ability of an atom in a molecule to attract electrons to itself (in a bond) F = 4.0 • EN differences: Covalent < 1.67 < Ionic

36. Bond Polarity and Electronegativity Assign a positive and negative pole to each of the following pair, decide which is ore polar B - F B - Cl Si - O P - P C = O C = S

37. Combining Formal Charge and Bond Polarity • Electroneutrality Principle electrons in a molecule are distributed in such a way that the charges on the atoms are as close to 0 as possible. • When a negative charge occurs, it is on the most electronegative element. • BF4- formal charge gives B a (-) but actually the (-) is distributed over the F

38. Combining Formal Charge and Bond Polarity A is more satisfactory

39. Combining Formal Charge and Bond Polarity

40. Bond Properties- Bond Order • The Order of a Bond is the number of bonding electron pair shared by two atoms. • Bond order = 1single H2 CH4 • Bond order = 2 double H2 C = C H2 • Bond Order = 3 triple H C C H

41. Bond Properties - Bond Order for Resonance Structures Bond Order = number of shared pair number of bonding sites

42. Bond Length and Bond Order Bond Length is the distance Between the nuclei of two Bonded atoms. Single bonds are determined By the size of atoms Double bonds are shorter Triple bonds, shorter still.

43. Bond Length and Bond Order

44. Bond Length and Bond Order

45. Bond Energy: Bond Dissociation Energy

46. Bond Energy: Bond Dissociation Energy Bonds broken C=C 610 kJ / mol and the H-H 1046 kJ / mol Bonds formed C-C 346 kJ and 2 C-H 2 x 413 kJ / mol D Horxn = 1046 kJ - 1172 kJ = -126 kJ

47. Molecular Shapes

48. Molecular Shapes • VSEPR valence shell electron-pair repulsion theory • Lewis dot diagrams can predict the 3-d geometry of a molecule - which determines its properties • Bond and lone electron pair in the valence shell of an element repel each other and seek to be as far apart as possible. • Central Atoms Surrounded Only by Single -Bond Pairs • Linear - trigonal planar -Tetrahedral • Trigonal-bipyrimadal -octahedral

49. Molecular Shapes single bond -Pair

50. Molecular Shapes single bond -Pair Expanded Octet