Chemical Bonding I: Lewis Theory

1. Chemical Bonding I: Lewis Theory Chapter 9

2. Chemical Bonding Atoms gain, lose, or share electrons in order to achieve a full outer shell electron configuration.

3. Ionic Bonds Composed of ions that have gained or lost electrons to achieve a full outer shell Electrostatic attractive forces Crystalline solids – no discrete molecules - formula units Identified by empirical formulas Metal + non-metal Covalent Bonds Composed of atoms that are sharing electrons to achieve a full outer shell Shared electron bonds Discrete molecules, forms gases, liquids, and solids Identified by molecular formulas Non-metal + non-metal Bonds

4. Molecular compounds boil at low temperatures because only weak intermolecular forces must be disrupted.

5. Ionic compounds boil at high temperatures because strong electrostatic bonding forces must be disrupted.

6. Ionic Bonds • Ions are held together by ionic attraction where the force of attraction is governed by Coulomb’s Law. • Makes sense • Large Z  strong attraction  larger E • Large d  charge felt less  smaller E Z=ion charge d = distance between nuclei

8. Lattice Energies

9. Lattice energy dependence on atomic sizes

10. Lattice energy dependence on ionic charges

11. Formation of ionic compounds • Get elements as atoms (generally requires energy) • Form ions (anions are energetically favorable, cations are unfavorable) • Bring ions together (favorable) • Condense to solid phase (favorable)

12. Energetics of NaCl Formation • Na(s) Na(g) +107.3 kJ/mol • Na(g)  Na+ + 1e +495.8 kJ/mol • 1/2 Cl2(g)  Cl (g) +122 kJ/mol • Cl(g)  Cl(g) 348.6 kJ/mol • Na+(g) + Cl(g)  NaCl(s) 787 kJ/mol • ==================================== • Na(s) + 1/2 Cl2(g)  NaCl(s) 411 kJ/mol

15. Determine the energy of formation of MgBr2 from the elements.

19. magnesium bromide • Mg(s)  Mg(g) +147.7 kJ/mol • Mg(g)  Mg+ + 1e +737.7 kJ/mol • Mg+(g)  Mg2+(g) + 1e +1450.7 kJ/mol • Br2(g)  2 Br (g) +193 kJ/mol • 2Br(g)  2Br(g) 2(325 kJ/mol) • = 650 kJ/mol • Mg2+(g) + 2Br(g)  MgBr2(s) 2440 kJ/mol • ==================================== • Mg(s) + Br2(g)  MgBr2(s)  561 kJ/mol

20. Mg + Cl2 MgCl2DH = -642 kJ/molMg + Br2  MgBr2DH = -561 kJ/molWhy are they different?

21. Calculate the energy released in kJ/mol in the reaction • Na(s) + 1/2 I2(s)  NaI(s) • The energy of vaporization of Na(s) is 107 kJ/mol. The sum of the • enthalpies of dissociation and vaporization of I2(s) is 214 kJ/mol, and the lattice energy of NaI is 704 kJ/mol.

22. Calculate the energy released in kJ/mol when LiH is formed in the reaction • Li(s) + ½ H2(g)  LiH(s) • Heat of vaporization, Li 161 kJ/mol • Dissociation energy, H2 436 kJ/mol • Lattice energy, LiH -917 kJ/mol • Ionization energy, Li 520 kJ/mol • Electron affinity, H -73 kJ/mol • Answer: -91 kJ/mole net change

23. Covalent Bonding • Shared electron bonds • Due to overlap of atomic orbitals • (Valence Bond Theory) • Allows each atom to fill valence shell with electrons

24. Representing Atoms, Ions, and Molecules as Lewis Electron Dot Structures • Use dots to represent valence electrons

26. Polar Covalent Bonds and Electronegativity • Polar bonds are bonds where the electron density is not shared equally between the two bonded atoms. • In polar bonds there is a positive and a negative end to the bond.

28. Bond Polarity NaCl HCl Cl-Cl

29. Electronegativity • The ability of an atom in a bond to attract electrons toward itself. • Electron greed • Note that electronegativity increases up and to the right as do the ionization energy and the electron affinity

31. Lewis Electron Dot Structures • Bonding electrons pairs – electron pairs involved in bonds • Lone electron pairs – electron pairs that do not participate in bonding • Bond order = number of bonds

32. Writing Lewis Dot Structures • Decide which atoms are bonded together - draw a skeleton structure • Count the total number of valence electrons available. • Find the number of electrons needed to give an octet around all atoms -- (remember H needs 2, all else need 8).

33. Writing Lewis Dot Structures • Determine number of electrons short. • Number of bonds needed = number of electrons short/2. • Distribute bonds -- (1st hook atoms together and then add double bonds where appropriate). • Calculate number of electrons used in bonds.

34. Writing Lewis Dot Structures • Calculate electrons remaining. • Distribute remaining electrons to give all atoms an octet. • Done!!

35. Formal Charge • The result of a method of electron bookkeeping that tells whether an atom in a molecule has gained or lost electrons compared to an isolated atom. • Formal charge = # valence electrons – (# bonds + # electrons as lone pairs)

36. Expanded Octets • Elements beyond neon have available d orbitals that may be used to accept additional electrons if necessary. • If you the number of bonds necessary to hook all atoms together is greater than the number needed to give all an octet then put in necessary bonds and distribute extra electrons on atoms that have available d orbitals in which to expand.

37. Lewis Structures of ions • for anions add the extra electrons to the number available • for cations subtract the lost electrons from the number available

38. Resonance • In some Lewis structures, the multiple bonds can be written in several equivalent locations. All structures have the exact same energy. Which is the correct Lewis structure?? • Answer : None alone are correct – the true molecule is a hybrid of the possible structures. The electrons are delocalized.

39. Bond length - the optimum distance between nuclei in a covalent bond.

40. Bond dissociation energy – (Bond Strength) • the amount of energy necessary to break a chemical bond in an isolated molecule in the gaseous state • the amount of energy released when a bond forms • Average bond dissociation energies are tabulated in the book. • Table 9.3

41. It is experimentally found that there is a direct correlation between the bond length and the bond strength. As the bond length decreases, the bond strength increases.

42. Sigma bond the first bond to form between any two atoms forms between atoms Pi bond second or third bond to form between two atoms forms above and below plane of the molecule Definitions