Chapter 6. Bonding

1. Chapter 6. Bonding 6.1 Types of Chemical Bonds 6.2 Electronegativity 6.3 Bond Polarity and Dipole Moments 6.4 Ions: Electron Configurations and Sizes 6.5 Formation of Binary Ionic Compounds 6.6 Partial Ionic Character of Covalent Bonds 6.7 The Covalent Chemical Bond: A Model 6.8 Covalent Bond Energies and Chemical Reactions 6.9 The Localized Electron Bonding Model 6.10 Lewis Structure 6.11 Resonance 6.12 Exceptions to the Octet Rule 6.13 Molecular Structure: The VSEPR Model

2. What is a Chemical Bond? (1) • Bonding is the force of attraction that holds atoms together in an element (N2) or compound (CO2 or NaCl). • The distances between bonded atoms are less than those between non-bonded atoms. • The forces between bonded atoms are greater than those between non-bonded atoms. • The principal types of bonding are ionic, covalent, and metallic.

3. What is a Chemical Bond? (2) • A chemical bond links two atoms or groups of atoms when the forces acting between them are sufficient to lead to the formation of an aggregate (a molecule) with sufficient stability to make it convenient for the chemist to consider it as an independent "molecular species” Paraphrased from Linus Pauling (1967).

4. . . . . . . . . - - : Cl : : Cl : Types of Chemical Bonds (1) • Ionic Bonds • Ionic substances are formed when an atom that loses electrons easily reacts with an atom that gains electrons easily. Loss of an electron Na · → Na++ e- e-+ Cl → Gain of an electron Combination to form the compound NaCl NaCl + Na+ →

5. For ionic bonds, the energy of interaction between a pair of ions can be calculated by using Coulomb's law. The energy depends only on distance. Types of Chemical Bonds (2)

6. Types of Chemical Bonds (3) • Covalent Bond • is the sharing of pairs of electrons between atoms. Covalent bonding does not require atoms be the same elements but that they be of comparable electronegativity. Covalent bonds give an the angular relation between the atoms (in polyatomic molecules, does not apply to molecules like H2). H· ·H → H:H or H―H

7. the tendency of an atom in a molecule to attract shared electrons. Shared electrons are closer to atoms with greater electronegativity. Trends in EN In a group (column): EN decreases w/ increasing Z (# of protons) In a period (row): EN increases w/ increasing Z Zumdahl Chapter 13 Electronegativity (EN)

8. The effect of an electric field on hydrogen fluoride molecules. Polar covalent bonds Dipole Moments • Bonded atoms share electrons unequally, whenever they differ in Electronegativity • HF: The F atom carries negative charge and the H atom positive charge of equal magnitude. The molecules align in an electric field. • Polar molecules posses a dipole moment, μ

9. Covalent/Polar/Ionic Percent Ionic Character Covalent e.g., H2, Cl2 N2 Polar Covalent e.g., HF, H2O Ionic Bond e.g., LiF, NaCl

10. Ionic vs. Covalent Bonding

11. The Pauling Electronegativity values as updated by A.L. Allred in 1961. Electronegativity

12. The Process of Bond Formation (1). • Two atoms with (i) unfilled electron shells or (ii) opposite charge, are separated by a great distance.The initial potential energy is zero. The atoms do not sense each other. • Via some random process the atoms approach each other.There is an attractive force between the atoms. The PE goes negative. • The distance between the atoms reaches the ideal bond distance.The PE reaches a minimum, attraction equals the repulsion. There is no net force. • The distance continues to decreaseRepulsion goes up sharply. PE goes positive. The atoms are forced back to the ideal bond distance.

13. The Process of Bond Formation (2).

14. The Process of Bond Formation (3) • When two atoms (or molecules) approach each other, the result depends on the atom (or molecule) type. • Cl + Cl both have unfilled shells. A bond can form. • K+ + Cl- have opposing charges. A bond can form. • Ar + Ar both have filled shells. A bond cannot form.

15. Electrons can be divided into: Valence electrons (e- in unfilled shells, outermost electrons). Valence electrons participate in bonding. Core electrons (e- in a filled shells). Core electrons do not participate in bonding. Valence electrons in molecules are usually distributed in such a way that each main-group elementis surrounded byeight electrons (an octet of electrons). Hydrogen is surrounded by two valence electrons. Core Electrons vs Valence Electrons

16. Sizes of Ions Cations (+): smaller than parent atom Anions (-): larger than parent atom Isoelectronic series: O2- F- Na+ Mg2+ Al3+ As Z, the nuclear charge, increases the atomic radii decreases Ionic Radii In picometers

17. Ions: Predicting Forumlae of Salts K: [Ar]4s1 Na+: [Ar] Can lose 1 electron Ca: [Ar]4s2 Ca+2: [Ar] Can lose 2 electrons O-2: [Ne] O: [He] 2s22p4 Can gain 2 electrons

18. He Dot Structures: closed shell elements 2 8 8 2 8 Ne Ar 2 8 8 18 Kr 2 Xe2, 8, 8, 18, 18 Rn 2, 8, 8, 18, 18, 32

19. Dot Structures: steps in converting a molecular formula into a Lewis Dot structure Draw the individual atoms/ions - with their valence electrons. H has 1, Carbon has 4, N has 5, O has 6, etc. Determine the total number of valence electrons. Account for the net charge. Step 1 Molecular formula Make the molecule. Place the atom with lowest EN in center (CO2, SO4 ). Step 2 Atom placement Make the bonds. Add 2 e- to each bond Step 3 Follow the octet rule. Add remaining e- (double, triple bonds, account for net charge) Make single bonds Step 4 Impose octet rule to C, N, O, F, Check the formal and total charges, and octet rule. Step 4 Lewis structure

20. Lewis Dot Structure NF3 Example NF3

21. Lewis Dot Structure: NF3

22. Dot Structures: NF3 Molecular formula count valence e- Atom placement Add in valence e- done

23. Dot Structures PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone. Write Lewis Structures for Molecules with One Central Atom

24. Dot Structures PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone. : : : Cl : : : Cl : C F : : : : F : The Lewis Structures for Molecules with One Central Atom SOLUTION:

25. Dot Structures PROBLEM: Write the Lewis structure for methanol (molecular formula CH4O, i.e., CH3OH), an important industrial alcohol that is used as a gasoline alternative in car engines. If you drink it you get drunk and then go blind and die. The Lewis Structure for Molecules with More than One Central Atom Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds.

26. Dot Structures PROBLEM: Write the Lewis structure for methanol (molecular formula CH4O, i.e., CH3OH), an important industrial alcohol that is used as a gasoline alternative in car engines. If you drink it you get drunk and then go blind and die. The Lewis Structure for Molecules with More than One Central Atom Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds. There are 4(1) + 4 + 6 = 14 valence e-. C forms 4 bonds, O forms 2. Each H forms 1 bond. O has 2 paisr of nonbonding e-. H : H C O H : H

27. Dot Structures Writing Lewis Structures for Molecules with Multiple Bonds. PROBLEM: Write Lewis structures for the following: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers. (b) Nitrogen (N2), the most abundant atmospheric gas. Try O2, H2.

28. Dot Structures H H H H C C C C H H H H The Lewis Structures for Molecules with double or triple bond. PROBLEM: Write Lewis structures for: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas If a central atom does not have 8e-, an octet, then a pair of e- can be moved from an adjacent atom to form a multiple bond. PLAN: (a) There are 2(4) + 4(1) = 12 valence e-. H can can form only one bond. Ethylene SOLUTION: :

29. Dot Structures N : N : . . N : N : N : N : . . . . The Lewis Structures for Molecules with double or triple bond. PROBLEM: Write Lewis structures for: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas If a central atom does not have 8e-, an octet, then a pair of e- can be moved from an adjacent atom to form a multiple bond. PLAN: N2 SOLUTION: (b) N2 has 2(5) = 10 valence e-. Therefore a triple bond is required to make the octet around each N.

30. Zumdahl Chapter 13 Dot Structures Write the Lewis Structure for O3 (Ozone)

31. Resonance is used to indicate that resonance occurs. Delocalized Electron-Pair Bonding O3 can be drawn in 2 ways - Neither structure is acurate. Reality is hybrid of the two. Resonance structures have the same relative atom placement but a difference in the locations of bonding and nonbonding electron pairs.

32. Formal Charge For OC For OA # valence e- = 6 # valence e- = 6 # nonbonding e- = 6 # nonbonding e- = 4 # bonding e- = 2 X 1/2 = 1 # bonding e- = 4 X 1/2 = 2 For OB Formal charge = -1 Formal charge = 0 # valence e- = 6 # nonbonding e- = 2 # bonding e- = 6 X 1/2 = 3 Formal charge = +1 Selecting the Best Resonance Structure An atom “owns” all of its nonbonding electrons and half of its bonding electrons. Formal charge is the charge an atom would have if the bonding electrons were shared equally. Formal charge of atom = # valence e- - (# unshared electrons + 1/2 # shared electrons)

33. Resonance SAMPLE PROBLEM 10.4 Writing Resonance Structures Write resonance structures for the nitrate ion, NO3-. PROBLEM:

34. Resonance SAMPLE PROBLEM 10.4 Writing Resonance Structures Write resonance structures for the nitrate ion, NO3-. PROBLEM: Use 5+(3x6)+1=24 valence e-. After you write out the dot structure see if other structures can be drawn in which the electrons can be delocalized over more than two atoms. PLAN: SOLUTION: Nitrate has 1(5) + 3(6) + 1 = 24 valence e- Something is wrong. N does not have an octet; shift a pair of e- to form a double bond

35. Resonance None of these are the ‘true’ structure. The true structure is a average of all three, and cannot be represented by a simple dot structure.

36. Resonance and Formal Charge Problem: Draw all the resonance structures of NCO-. Determine the relative weights of the resonance structures.

37. Resonance and Formal Charge A B C Problem: Draw all the resonance structures of NCO-. formal charges -2 0 +1 -1 0 0 0 0 -1 The true structure is a weighted average of Forms A, B and C. Form A: Formal charge of -2 on N and +1 on O. O is the most electro-negative atom in the molecule. Low weight. Form B: Formal charge on the N is more negative than that of O. O is the most electro-negative atom in the molecule. Low weight. Form C: Formal charge on O is -1. High weight.

38. Resonance and Formal Charge Three criteria for estimating the weights of resonance structures: Smaller formal charges (absolute value) give higher weight than larger charges. Opposing charges on adjacent atoms gives a low weight. A negative formal charge is on an electronegative atom (O and sometimes N and S) gives a high weight.

39. Resonance and Weighted Averages A B C EXAMPLE: NCO- has 3 possible resonance forms - None of these are the ‘true’ structure. The true structure is a weighted average of all three. The weight of C is greater than the weights of A and B.

40. Carbon Monoxide

41. Dot Structures - Octet Expansion & Contraction PROBLEM: Write Lewis structures for (a) H3PO4 (pick the most heavily weighted structure); (b) BFCl2. PLAN: Draw the Lewis structures for the molecule. Note that these dare exceptiosn to the octet rule. P is a Period-3 element and can have an expanded valence shell. SOLUTION: (a) H3PO4 has two resonance forms and formal charges indicate the more important form. -1 0 +1 0 (b) BFCl2 has only 1 Lewis structure. 0 0 0 0 0 0 0 0 0 0 more weight 0 0 lower formal charges

42. VSEPR - Valence Shell Electron Pair Repulsion Theory Each electron pair (bonded and non-bonded) around a central atom is located as far away as possible from the others, to minimize electrostatic repulsions. But a double bond counts as one bond. A triple bond counts as one bond. These repulsions maximize the space allowed for each electron pair. The result is five electron-group arrangements of minimum energy seen in a large majority of molecules and polyatomic ions. Linear, trigonal planer, tetrahedral, trigonal bipyramidal, octahedral.

43. Electron-pair repulsions and the five basic molecular shapes. linear tetrahedral trigonal planar trigonal bipyramidal octahedral

44. The single molecular shape of the linear electron-group arrangement. Examples: CS2, HCN, BeF2

45. Class Shape Figure 10.4 The two molecular shapes of the trigonal planar electron-group arrangement. Examples: SO2, O3, PbCl2, SnBr2 Examples: SO3, BF3, NO3-, CO32-

46. Effect of Double Bonds Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 larger EN 1200 ideal greater electron density

47. 1220 Effect of Double Bonds 1160 real Effect of Nonbonding(Lone) Pairs Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 larger EN 1200 ideal greater electron density Lone pairs take more space than bonding pairs.. 950

48. 1220 Effect of Double Bonds 1160 real Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 larger EN 1200 ideal greater electron density

49. 1220 Effect of Double Bonds 1160 real Effect of Nonbonding(Lone) Pairs Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 larger EN 1200 ideal greater electron density Lone pairs repel bonding pairs more strongly than bonding pairs repel each other. 950

50. The three molecular shapes of the tetrahedral electron-group arrangement. Figure 10.5 Examples: CH4, SiCl4, SO42-, ClO4- NH3 PF3 ClO3 H3O+ H2O OF2 SCl2