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. 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.
Na · →
Na+ + e-
Loss of an electron
: Cl :
. . -
e- + Cl →
Gain of an electron
: Cl :
. . -
Combination to form the compound NaCl
NaCl
Na+ + →
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5. Types of Chemical Bonds (2)
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.

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
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7. Electronegativity (EN)
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
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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. Electronegativity
The Pauling Electronegativity values as updated by A.L. Allred in 1961.

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 decrease Repulsion 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. Core Electrons vs Valence Electrons
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 element is surrounded by eight electrons (an octet of electrons). Hydrogen is surrounded by two 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. Sizes of ions

18. 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

19. Dot Structures: closed shell elements2
Xe 2, 8, 8, 18, 18 Ne 2 8 Ar 2 8
Rn 2, 8, 8, 18, 18, 32 Kr 2 8 18

20. Dot Structures: valence electrons and the periodic table

21. 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
Step 3
Make the bonds. Add 2 e- to each bond
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

22. Lewis Dot Structure NF3 Example

23. Lewis Dot Structure: NF3

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

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

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

27. Dot Structures
The Lewis Structure for Molecules with More than One Central Atom
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.
Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds.

28. Dot Structures H : H C O H : H
The Lewis Structure for Molecules with More than One Central Atom
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.
Hydrogen can have only one bond. C and O are bonded. H fills in the rest of the bonds.
There are 4(1) = 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

29. Dot Structures Write Lewis structures for the following:
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.

30. Dot Structures
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:
Ethylene SOLUTION:
(a) There are 2(4) + 4(1) = 12 valence e-. H can can form only one bond. C H C H :

31. Dot Structures
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. N : . N : . N :

32. Write the Lewis Structure for O3 (Ozone)
Dot Structures
Write the Lewis Structure for O3 (Ozone)
Zumdahl Chapter 13
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33. Resonance 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.
is used to indicate that resonance occurs.

34. Formal Charge 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)
For OC
# valence e- = 6
# nonbonding e- = 6
# bonding e- = 2 X 1/2 = 1
Formal charge = -1
For OA
# valence e- = 6
# nonbonding e- = 4
# bonding e- = 4 X 1/2 = 2
Formal charge = 0
For OB
# valence e- = 6
# nonbonding e- = 2
# bonding e- = 6 X 1/2 = 3
Formal charge = +1

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

36. 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

37. 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. 37

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

39. Resonance and Formal Charge
Problem: Draw all the resonance structures of NCO-. A B C
formal charges -2 +1 -1 -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.

40. 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.

41. Resonance and Weighted Averages
EXAMPLE: NCO- has 3 possible resonance forms - A B C
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. 41

42. Carbon Monoxide

43. 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 +1
(b) BFCl2 has only 1 Lewis structure.
more weight
lower formal charges

44. 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.

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

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

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

48. 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
Effect of Double Bonds
larger EN
1200
ideal
greater electron density

49. 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
1220
1160
real
Effect of Double Bonds
larger EN
1200
ideal
greater electron density
Effect of Nonbonding(Lone) Pairs
Lone pairs take more space than bonding pairs..
950

50. 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
1220
1160
real
Effect of Double Bonds
larger EN
1200
ideal
greater electron density

51. 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
1220
1160
real
Effect of Double Bonds
larger EN
1200
ideal
greater electron density
Effect of Nonbonding(Lone) Pairs
Lone pairs repel bonding pairs more strongly than bonding pairs repel each other.
950

52. 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

53. Lewis structures and molecular shapes.
Figure 10.6
Lewis structures and molecular shapes.

54. Figure 10.7
The four molecular shapes of the trigonal bipyramidal electron-group arrangement.
PF5
AsF5
SOF4
SF4
XeO2F2
IF4+
IO2F2-
XeF2
I3-
IF2-
ClF3
BrF3

55. Figure 10.8
The three molecular shapes of the octahedral electron-group arrangement.
SF6
IOF5
BrF5
TeF5-
XeOF4
XeF4
ICl4-

56. A summary of common molecular shapes with two to six electron groups.

57. SAMPLE PROBLEM
Predicting Molecular Shapes with Five or Six Electron Groups
PROBLEM:
Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5.

58. SAMPLE PROBLEM 10.7
Predicting Molecular Shapes with Five or Six Electron Groups
PROBLEM:
Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5.
SOLUTION:
(a) SbF valence e-; all electrons around central atom will be in bonding pairs; shape is AX5 - trigonal bipyramidal.
(b) BrF valence e-; 5 bonding pairs and 1 nonbonding pair on central atom. Shape is AX5E, square pyramidal.

59. SAMPLE PROBLEM
Predicting Molecular Shapes with More Than One Central Atom
PROBLEM:
Determine the shape around each of the central atoms in acetone, (CH3)2C=O.
PLAN:
Find the shape of one atom at a time after writing the Lewis structure.
SOLUTION:
tetrahedral
tetrahedral
trigonal planar
>1200
<1200

60. The tetrahedral centers of ethane and ethanol.
Figure 10.11
The tetrahedral centers of ethane and ethanol.
ethanol
CH3CH2OH
ethane
CH3CH3

61. The orientation of polar molecules in an electric field.
Figure 10.12
The orientation of polar molecules in an electric field.
Electric field ON
Electric field OFF

62. Dipole moment = μ = QR
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63. Non-Polar covalent bonding
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64. SAMPLE PROBLEM 10.9
Predicting the Polarity of Molecules
PROBLEM:
From electronegativity (EN) values (button) and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole when applicable:
(a) Ammonia, NH3
(b) Boron trifluoride, BF3
(c) Carbonyl sulfide, COS (atom sequence SCO)

65. SAMPLE PROBLEM
Predicting the Polarity of Molecules
continued
(b) BF3 has 24 valence e- and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar.
F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar.
1200
(c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar(DEN) so the molecule is polar overall.

66. SAMPLE PROBLEM
Predicting the Polarity of Molecules
continued
(b) BF3 has 24 valence e- and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar.
F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar.
1200
(c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar(DEN) so the molecule is polar overall.

67. Carbonate, CO3-2
Thiocyanate ion, NCS-1
N2O, laughing gas

68. SN = 5 (bonded atoms) + 0 (lone pairs) 5: Trigonal bipyramidal
The VSEPR Theory
SN = 5 (bonded atoms) + 0 (lone pairs)
5: Trigonal bipyramidal
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69. SN = 6 (bonded atoms) + 0 (lone pairs) 6: Octahedral
The VSEPR Theory
SN = 6 (bonded atoms) + 0 (lone pairs)
6: Octahedral
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70. Pyramidal The VSEPR Theory SN = 3 (bonded atoms) + 1 (lone pairs)
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71. Bent SN = 2 (bonded atoms) + 2 (lone pairs) Zumdahl Chapter 13
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72. Zumdahl Chapter 13
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73. Zumdahl Chapter 13
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74. Zumdahl Chapter 13
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75. Zumdahl Chapter 13
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