Chapter 9—Models of Chemical Bonding

1. Chapter 9—Models of Chemical Bonding AP Chemistry

2. What is a chemical bond? Forces that cause a group of atoms to behave as a unit

3. Why do atoms bond? To lower the potential energy of their atoms which creates a more stable arrangement of atoms

4. Figure 9.1 A general comparison of metals and nonmetals

5. Types of Chemical Bonding 1. Metal with nonmetal: electron transfer and ionic bonding 2. Nonmetal with nonmetal: electron sharing and covalent bonding 3. Metal with metal: electron pooling and metallic bonding

6. Figure 9.2 The three models of chemical bonding

7. Ionic Bond An atom that loses electron(s) combines with an atom that gains electron(s) Ex: Draw electron configurations for… Na Cl 1s22s22p6 1s22s22p63s23p Na+ Cl- 1 6 5 3s electrostatic attraction

8. Covalent Bond • Sharing of electron pairs between two atoms • Shared electrons are simultaneously attracted to both atomic nuclei Video

9. Structure of a metallic crystal The “sea” of delocalized electrons allows atoms to move past one another as the solid deforms, creating a malleable substance.

10. Gilbert Newton Lewis (1875-1946) • Determined electrode potentials, conductivity, free energy and other thermodynamic constants for elements. • Provided the first description of covalent bonding and a shared pair of electrons between atoms. Theorized that in most atoms electrons arranged themselves so that there were 8 electrons around the atoms. • First to prepare pure deuterium; predicted presence of heavy water. • Lewis's speculations onatomic structure.

11. The A group number gives the number of valence electrons. Place one dot per valence electron on each of the four sides of the element symbol. Pair the dots (electrons) until all of the valence electrons are used. . . . : . . . : N . N . . N N : . . : . Lewis Electron-Dot Symbols For main group elements - Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.

12. Dot Diagrams outer shell electrons • Symbolize (s and p orbitals only) • Why? Those are the electrons involved in bonding and chemical reactions LEFTRIGHTTOPBOTTOM X

13. Figure 9.3 Lewis electron-dot symbols for elements in Periods 2 and 3

14. PROBLEM: Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound. O2- Na Na 2s 2p 3s 3p O 2s 2p . . Na : 3s 3p + O : : 2Na+ + O 2- : : . : . Na SAMPLE PROBLEM 9.1 Depicting Ion Formation PLAN: Draw orbital diagrams for the atoms and then move electrons to make filled outer levels. It can be seen that 2 sodiums are needed for each oxygen. SOLUTION: 2 Na+

15. Octet Rule • When atoms bond, they lose, gain, or share electrons to attain a filled outer level of eight (or two) electrons

16. Ionic Bonding • Transfer of electrons from metal atoms to nonmetal atoms to form ions that come together in a solid ionic compound

17. + F- 1s22s22p6 Na+ Na 3s 3p 3s 3p F- + + F 1s 2s 2p 1s 2s 2p . : . : F : Na+ + : F - : Na + : : Three ways to represent the formation of Na+ and F- through electron transfer. Figure 9.4 Electron configurations Na 1s22s22p63s1 + F 1s22s22p5 Na+ 1s22s22p6 Orbital diagrams Lewis electron-dot symbols

18. Bond Energy Energyrequiredtobreaka bond (kJ/mol)(+) Energyreleasedwhen a bond isformed(-)

19. 4 3 5 2 1 Overall Energy Change 6 Calculate DE In Forming NaF 12: Na(s)Na(g)DE = +109 kJ energy of sublimation 23: Na(g) Na+ (g) DE = +495 kJ Ionization Energy 34: ½ F2 (g)  F(g) DE = +77 kJ E to break F-F bond 45: F(g) + 1e- F-1DE = -328 kJ Electron Affinity 56: Na+(g) + F-(g)  NaF(g) DE = -923 kJ Lattice Energy DEnet = -570 kJ/mol

20. Figure 9.6 The Born-Haber cycle for lithium fluoride

21. Which is bigger? The formation of MgO or NaF? Why is Oxygen’s second electron affinity endothermic?

22. MgO Mg  Mg2+ + 2e- DH = 2188 kJ O + 1e- O-1DH = -141 kJ O-1 + 1e-  O-2 DH = 878 kJ O + 2e-  O-2DH = 737 kJ Why does MgO form???????? Mg2+ + O-2 MgO DH = 3923 kJ/mol

23. Lattice Energy The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid Ex: M+ (g) + X-(g) MX(l) + Energy Q is charge of ion R is distance b/w ions LE depends on charge more than radius/size

25. Sizes of Ions On P.T.

26. As size of ions increase, Lattice energy ________ As size of charge increase, Lattice energy ________

27. Sample Problem Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. NaCl vs. KCl MgO vs. NaCl Fe(OH)2 vs Fe(OH)3

28. NaCl vs. KCl The larger the ionic radius, the lower the lattice energy The amount of interaction between the ions is smaller and the packing of the ions is less efficient L.E. NaCl = -780 kJ/mol L.E. KCl = -711 kJ/mol

29. MgO vs. NaCl The stronger the charge on an ion the stronger the attractive force that will result in an ionic lattice. +/-1 ions form compounds with lower lattice energies than +/-2 ions. L.E. MgO = -3791 kJ/mol L. E. NaCl = -780 kJ/mol

30. Fe(OH)2 vs. Fe(OH)3

32. Solid ionic compound Molten ionic compound Ionic compound dissolved in water Figure 9.9 Electrical Conductance and Ion Mobility

33. Summary of Ionic Bonds Between metal ion and non – metal ion Electrostatic attraction between ions Veeerrrrrryyyyy strong!

34. Covalent Bond • Sharing of electron pairs between two atoms • Shared electrons are simultaneously attracted to both atomic nuclei

35. a d b c Relationship b/w Bond Length & Bond Energy Graph: Chang Disc 2 • Large distances away; no influence on each other; Ep = 0 • Atoms move closer together, nuclei are attracted to each other;Ep↓ • c) Attractive force dominates; Ep is at a minimum STABLE! • d) If nuclei move closer together, repulsive force > attractive force; Ep↑

37. Lone Pairs vs. Bonding Pairs Lone pairs: electron pairs not involved in bonding Bonding Pairs: electron pairs involved in bonding

38. Bond Length Distance between two atomic nuclei when energy is at its minimum (when they are the most stable)

39. Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Internuclear distance (bond length) Covalent radius Figure 9.13 Bond length and covalent radius.

40. Relationship between Bond Energy and Bond Length

44. (b) C = O, C - O, C O Bond length: C - O > C = O > C O Bond strength: C O > C = O > C - O SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength PROBLEM: Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength: (a) S - F, S - Br, S - Cl PLAN: (a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases. SOLUTION: (a) Atomic size increases going down a group. (b) Using bond orders we get Bond length: S - Br > S - Cl > S - F Bond strength: S - F > S - Cl > S - Br

45. Strong covalent bonding forces within molecules Weak intermolecular forces between molecules Strong forces within molecules and weak forces between them. Figure 9.14

46. Figure 9.15 Covalent bonds of network covalent solids.

47. Covalent Bond Energies & Chemical Reactions Recall relationship between bond length and bond energy…

48. Bond Energy – Average of Individual Bond Energies Bond energy depends on the environment

49. Types of Covalent Bonds single bond— double bond— triple bond— 1 e- pair shared 2 e- pairs shared 3 e- pairs shared

50. Bond Type