CHAPTER 9: BONDING AND MOLECULAR STRUCTURE

1. CHAPTER 9: BONDING AND MOLECULAR STRUCTURE

2. 9.0 Objectives: Understand the basic process of ionic bonding: identify ionic compounds and describe their internal structure and properties, understand how size of ions affects ionic properties, calculate lattice energy, and draw Lewis diagrams of ionic structures. Identify covalent compounds and characterize their properties. Draw Lewis structures of covalent substances including exceptions such as reduced and expanded octets, radicals, and resonance structures. Define and predict trends in bond order, bond length, and bond dissociation energy. Use bond energy to predict enthalpy of a reaction. Understand the concept of electronegativity and how it is used to predict polarity of individual bonds and entire molecules. Use VSEPR Theory to predict the shapes of simple covalent molecules.

3. Homework: • HW#1 – 29, 33, 38, 39, 40, 41, 43, 45 • Valence e-, LEWIS STRUCTURES! • HW#2 – 47, 49, 51, 53, 95 • Formal Charge, Polarity, Electonegativity • HW#3 – 69, 93, 109 • Bond Energy • HW#4 – 73, 75, 77, 79, 81, 89, 99, 103 • Molecular Geometry • HW#5 – 83, 85, 97 • Polarity

4. 9.1 VALENCE ELECTRONS • 1. Bonding – definition • “forces that hold atoms together” • What role do the e- play? • 2. Valence electrons vs. Core electrons • MAIN GROUP Elements • Outermost “s” and “p” e- • Transition Metals • Outermost “s” and “p” e- as well as (n-1) “d” e- When in doubt, write the noble gas e- configuration

5. 9.1 VALENCE ELECTRONS • 3. Lewis dot diagrams of elements • Diagrams that showcase valence e- • Lewis says, “Place the first four dots separately!” • Ex. Li, Be, B, C, N, O, F, Ne

6. 9.2 CHEMICAL BOND FORMATION • 1. Ionic bonding – definition and Lewis representation of NaCl • Bond between metal and nonmetal due to “electrostatic interactions” • Metal donates an e- • Nonmetal accepts an e- • Ex. NaCl

7. 9.2 CHEMICAL BOND FORMATION • 2. Covalent bonding and Lewis representation of Cl2 • Bond in which e- are shared • Overlap of e- density between 2 orbitals • Ex. Cl2

8. 9.2 CHEMICAL BOND FORMATION • 3. Continuum • Complete ionic or complete sharing of e- is a bit extreme; most bonding has uneven sharing of e- (sometimes ionic, sometimes covalent) • 4. Other bond types • Metallic bonding • Ex. Alloys

9. 9.3 BONDING IN IONIC COMPOUNDS • 1. Steps in formation of NaCl • 1. Na(g) Na+(g) + e- E = +496 kJ/mol • 2. Cl(g) + e-  Cl-(g)E = -349 kJ/mol • 3. Na+(g) + Cl-(g) [Na+, Cl-] E = -498 kJ/mol Eoverall = -351 kJ/mol

10. 9.3 BONDING IN IONIC COMPOUNDS • 2. Lattice energy • “energy for the formation of 1 mol of solid crystalline ionic compound when ions in the gas phase combine”

11. 9.3 BONDING IN IONIC COMPOUNDS • 3. Formula units • Smallest whole number Ratio -repeating unit of an ionic compound

12. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • 1. Diagram of H2 H H Both want 1s2 H:H Share the electron pair H—H Bonded stable H2

13. 9.4 COVALENT BONDING AND LEWIS STRUCTURES 2. Orbital overlap diagrams of H2, HCl, Cl2

14. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • 3. Terminology – single, double, and triple bonds, bonding pairs and nonbonding or lone pairs of electrons • Single Bond: 2 e- shared between 2 atoms • Ex. H2 • Double Bond: 4 e- shared between 2 atoms • Ex. O2 • Triple Bond: 6 e- shared between 2 atoms • Ex. N2

15. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • Bonding Pairs: e- involved in bonding • (See preceding examples) • Nonbonding (lone) pairs: e- that are not involved with bonding but help provide the octet for an atom • Ex. Cl2

16. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • 4. Octet Rule • Noble-gas configuration • “tendency for molecules/polyatomic ions to have structures in which 8 e- surround each atom” • H, He have a “duet” • Be has 4 electron max and B has 6 electron max

17. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • 5. Rules for drawing Lewis structures • a. Choose a central atom • Usually the atom with the lowest e- affinity • Usually makes a lot of bonds • Halogens are generally terminal atoms • b. Count the total number of valence electrons • Neutral Molecule: sum of valence e- for each atom • Anions: sum of valence e- and negative charge • Cations: valence e- minus the total positive charge

18. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • c. Draw a skeleton structure • Use one pair of electrons to form a bond between each pair of bound atoms • d. Place the remaining electrons to fulfill the octet rule • Do this for each atom • Hydrogen gets a duet

19. 9.4 COVALENT BONDING AND LEWIS STRUCTURES • e. Lack of electrons: • Requires multiple bonds (double, triple) • Could be more than one multiple bond • f. Too many electrons: • Verify that your structure is correct (octets for all?) • Watch anions!

20. 9.6 Lewis Structures of Some Simple Molecules • NOT suv….SOV! S = O-V S: Shared e- in bonds O: total # e- required for an Octet V: Valence e- for all elements

21. 9.4 COVALENT BONDING AND LEWIS STRUCTURES 6. Diagrams of H2 F2 CH4 NH3 H2O HF OH- NH4+

22. 9.4 COVALENT BONDING AND LEWIS STRUCTURES H2 F2 CH4 NH3 H2O HF OH- NH4+

23. 9.4 COVALENT BONDING AND LEWIS STRUCTURES 7. Isoelectronic species: NO+ N2 CO CN-

24. 9.5 RESONANCE • 1. Definition • Alternative and equivalent Lewis structure “created” by shifting the e- in a structure • Spinning Rim Analogy

25. 9.5 RESONANCE 2. Examples: NO3- and NO2-

26. 9.5 RESONANCE • 3. Experimental evidence says: • “It’s a combination of both” • There are however, MORE PREVALENT resonance structures for some molecules • Benzene is the most classic of all resonance structures

27. 9.5 RESONANCE

28. 9.6 EXCEPTIONS TO THE OCTET RULE • 1. Reduced octets for H, B and Be • Ex. BeCl2, BCl3 (Be = 4 e-, B = 6e-)

29. 9.6 EXCEPTIONS TO THE OCTET RULE • 2. Expanded octets: PF5 SF6 ClF4- XeF2 • Watch these elements (and some others) for expanded octets: P, S, Cl, As, Se, Br, Kr, Xe

30. 9.6 EXCEPTIONS TO THE OCTET RULE • 3. Radicals (paramagnetic): NO and NO2 • Structure that has unpaired e- • Extremely Reactive • O2 ?!?!

31. 9.6 EXCEPTIONS TO THE OCTET RULE • 3. Problems with Lewis structures • Only show 2-D view life (chemistry) is 3-D • Works for most molecules, but not all • Doesn’t show how evenly/unevenly e- are being shared

32. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 1. Definitions: polar and nonpolar bonds • Nonpolar bonds: 2 e- in a bond are “evenly” shared between the 2 atoms • Polar bonds: 2 e- in a bond are unevenly shared; one atom is taking more of the e- density; atoms have a partial charge

33. POLARITY RULES

34. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 2. Electronegativity • a. definition • (e/n): “ability of an atom to attract bonding e- to itself when the atom is in a molecule” • b. Table and Periodic trends • See Pg.10 in Reference Booklet • Increases going left to right and bottom to top • (Points toward Fluorine) • ACS Periodic Table

35. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES

36. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 3. EN – parameters • Prediction of “Ionic Character” Pure Covalent Pure Ionic 0 .5 1 1.5 2 2.5 3 In General: 0.0 < 0.5 Nonpolar 0.5 ≤ 1.8 Polar Covalent > 1.8 Ionic

37. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES 4. Ex9.1 Arrange the following bonds in order of increasing polarity: F-Cl, F-F, F-Na

38. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 5. Central atom in Lewis structure: • Many times has a formal charge • Making more/less bonds than it “normally” does

39. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 6. Formal Charge • a. Definition and Use • Charge for an atom in a molecule based on premise that bonding e- are evenly shared • b. Calculating – equation • Formal Charge = Group # - [Lone Pair e- + ½ Bonding e-]

40. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • c. Examples: OH- and NO3-

41. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES Ex9.2 Calculate the formal charge on each atom in CO32- and NH4+

42. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • 7. Electroneutrality – Definition • The e- in a molecule are distributed so that the formal charge is minimal • Most Probable Lewis Structure = one with minimal formal charge • Negative charge should reside on the most electronegative element • Formal charge > +/- 2 is not likely

43. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES • a. Example: CO2

44. 9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND MOLECULES b. Ex9.3 Use formal charge and the electroneutrality concept to determine the most likely structures for N2O and OCN1-

45. 9.8 BOND PROPERTIES • 1. Bond order • a. definition and examples • “number of bonding e- pairs shared between 2 atoms” • Usually an integer (1, 2, or 3) • BOND ORDER = (# shared pairs linking X-Y) (number of X-Y links in the molecule) • Ex. CH4, CO2 • b. resonance structures • Bond order are fractions • e- residing over both locations evenly • Ex. O3

46. 9.8 BOND PROPERTIES • 2. Bond Length – definition and examples • Bond length: distance between nuclei of a covalent bond (little variation) • More Polar bonds = shorter length • Ex. C-C C=C C≡C 1.54Å 1.34Å 1.20Å

47. 9.8 BOND PROPERTIES • 1. Bond dissociation energy – definition and examples • Bond Dissociation Energy (D): • “energy needed to break a covalent bond in the gas phase” • Higher Bond Order Higher Dissoc. Energy

48. 9.8 BOND PROPERTIES • a. Estimating Enthalpy of reaction from bond energies – equation • Hrxn =  D (bonds broken) -  D (bonds formed) • Energy is required to break bonds • Energy is released when bonds are formed

49. 9.8 BOND PROPERTIES

50. 9.8 BOND PROPERTIES • b. Example: Estimate the HR for the synthesis reaction between gaseous hydrogen and chlorine.