Chapter 13

1. Chapter 13 Bonding: General Concepts

2. Types of Chemical Bonds • Ionic bonding • Polar covalent bonding • Covalent bonding

4. Lennard-Jones 6-12 potential

5. Ionic bonding • Ionic substances are formed when an atom that loses electrons relatively easily react with an atom that has a high affinity for electrons. ex. metal-nonmetal compound

6. Covalent Bonding • Electron are shared by nuclei

7. Polar Covalent Bonding • A polar bond is a covalent bond in which there is a separation of charge between one end and the other , in other words in which one end is slightly positive and the other slightly negative.

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

9. Calculate Electronegativity • H atom ∆HF=565-1/2(432+154)=272 |xH-xF|=0.102(272)1/2=1.68, xH-4.0=-1.7, xH=2.3 • O atom ∆OF=190-1/2(146+154)=40 |xO-xF|=0.102(40)1/2=0.65, xO-4.0=-0.65, xO=3.4 • C atom ∆CF=485-1/2(347+154)=234.5 |xC-xF|=0.102(234.5)1/2=1.6, xC-4.0=-1.6, xC=2.4

13. Dipole Moment μ=QR Q: center of charge of magnitude R: distance Bond Polarity and Dipole Moments

14. Dipole Moment of HF 1D=3.336×10-30 coulomb meter μ=(1.6×10-19 C)(9.17×10-11 m)=1.47×10-29 =4.4 D for fully ionic Measured dipole moment=1.83 D 1.83×3.336×10-30=δ(9.17×10-11) δ=6.66×10-20 Ionic character=1.83/4.4=41.6%

15. In practice no bond is totally ionic. There will always be a small amount of electron sharing.

17. The compounds with more 50% ionic character are normally considered to be ionic solids.

18. Dipole Moment of Polyatomic Molecules • For dipole moment of polyatomic molecules, the dipole is the geometric sum of all bond dipole moment.

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

23. Isoelectronic Ions • Ions containing the the same number of electrons O2> F > Na+ > Mg2+ > Al3+ largest smallest

24. Formation of Binary Ionic Compounds Lattice energy: The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. M+(g)+X-(g) →MX(s)

25. 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 1/2X2(g) → X(g) [endothermic]

26. 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 M+(g) + X-(g) → MX(s) [quite exothermic]

27. Electron affinity of F Dissociation of F2 Ionization of Li Formation of solid Sublimation of Li

28. Lithium-Fluoride structure

29. Lattice Energy Calculations k: a proportionality constant that depends on the structure of the solid and the electron configuration of the ions Q1 and Q2: charges on the ions r: the shortest distance between the centers of cations and anions

31. Lattice Energies and the Strength of the Ionic Bond • The strength of the bond between the ions of opposite charge in an ionic compound depends on the charges on the ions and the distance between the centers of the ions when they pack to form a crystal. • An estimate of the strength of the bonds in an ionic compound can be obtained by measuring the lattice energy of the compound.

32. Lattice Energies for Alkali Metals Halides • The bond between ions of opposite charge is strongest when the ions are small. • The lattice energies for the alkali metal halides is therefore largest for LiF and smallest for CsI.

33. Lattice Energies of Alkali Metals Halides (kJ/mol)

34. Lattice Energies for Salts of the OH- and O2- Ions • The ionic bond should also become stronger as the charge on the ions becomes larger. • The lattice energies for salts of the OH- and O2- ions increase rapidly as the charge on the ion becomes larger.

35. Lattice Energies of Salts of the OH- and O2- Ions (kJ/mol)

36. Lattice Energies and Solubility • The lattice energy of a salt gives a rough indication of the solubility of the salt in water because it reflects the energy needed to separate the positive and negative ions in a salt. • Sodium and potassium salts are soluble in water because they have relatively small lattice energies. Magnesium and aluminum salts are often much less soluble because it takes more energy to separate the positive and negative ions in these salts. • NaOH is very soluble in water (420 g/L), but Mg(OH)2 dissolves in water only to the extent of 0.009 g/L, and Al(OH)3 is essentially insoluble in water.

37. The Covalent Chemical Bond

38. Bond Energy of CH4 Experimental result : 1652 kJ/mol C(g)+4H(g) →CH4(g) + 1652 kJ/mol An average C-H bond energy per mole of C-H bond: 1652/4=413 (kJ/mol)

39. Stepwise Decomposition of CH4 CH4(g) →CH3(g)+H(g) 435 kJ/mol CH3(g) →CH2(g)+H(g) 453 kJ/mol CH2(g) →CH(g)+H(g) 425 kJ/mol CH(g) →C (g)+H(g) 339 kJ/mol

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

41. Bond Energy of CH3Cl C(g)+Cl(g)+3H(g) →CH3Cl(g)+1578kJ/mol (C-Cl)+3(C-H)=1578 (C-Cl)+3(413)=1578 C-Cl=339 (kJ/mol)

42. Covalent Bond Energies and Chemical Reactions H2+F2→2HF ΔH=ΣD (bonds broken)-ΣD (bonds formed) ΔH=DH-H+DF-F-2DH-F=1×432+1×154-2×565 =-544 kJ

43. CH4+2Cl2+2F2→CF2Cl2+2HF+2HCl Reactants bonds broken: CH4: 4×413=1652, 2Cl2: 2×239=478 2F2:2×154=308 Total energy required: 2438kJ Products bonds formed: CF2Cl2: 2×485=970 (C-F) and 2×339=678 (C-Cl) HF: 2×565=1130 HCl: 2×427=854 Total energy released: 3632 kJ ΔH=2438-3632=-1194 kJ (-1126 kJ)

44. 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. • Description of valence electron arrangement (Lewis structure). • Prediction of geometry (VSEPR model). • Description of atomic orbital types used to share electrons or hold long pairs.

45. 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.

46. 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.

47. Comments About the Octet Rule • 2nd row elements C, N, O, F obey 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 CANexceed the octet rule using empty valence d orbitals. • When writing Lewis structures, satisfy octets first,then place electrons around elements having available d orbitals.

48. Formal Charge • One method involves estimating the charge on each atom in the various possible Lewis structures and using the charges to select the most appropriate structure. • It allows chemists to determine the location of charge in a molecule as well as compare how good a Lewis structure might be.

49. Calculation of Formal Charge • Formal Charge=(number of valence electrons on a free atom)-(number of valence electrons assigned to the atom in the molecule)

50. Assumptions for Formal Charge • Lone pair electrons belong entirely to the atom in question. • Shared electrons are divided equally between the two sharing atoms. (Valence electrons) assigned=(number of lone pair electrons)+1/2(number of shared electrons)