1. Chpt 8 - Bonding Covalent vs. Ionic Bond Polarity and dipole moments
(differences in electronegativity values)
Bond Polarity and dipole moments
Lattice Energy - ionic bonding
Covalent bond energy - molecules
Lewis Structures, resonance, exceptions
VSEPR
HW: Chpt 8 - pg , #s 15,19, 24, 28, 29, 30, 34, 38, 42, 44, 52, 53acd, 54acd, 61, 64, 70, 81, 82acd, 86, 87, 92, 97(look at 82) 102, 103, 108, 110, 118, Due Wed Nov. 10

2. What is Chemical Bonding?
What is meant by the term “chemical bond”?
Why do atoms bond with each other to form compounds?
How do atoms bond with each other to form compounds?

3. A Chemical Bond No simple, and yet complete, way to define this.
Forces that hold groups of atoms together and make them function as a unit.
A bond will form if the energy of the aggregate is lower than that of the separated atoms.

4. Energy Profile: H + H --> H2
The figure shows the interaction of 2 hydrogen atoms as the internuclear distance changes.

5. Types of Bonding Polar Covalent Bond
Ionic Bonding – electrons are transferred
Covalent Bonding – electrons are shared equally
What about intermediate cases?
Polar Covalent Bond
Unequal sharing of electrons between atoms in a molecule.
Results in a charge separation in the bond (partial positive and partial negative charge).

6. Electronegativity and Bond Type

7. Relative polarity of bonding with hydrogen

8. Polarity exercise a) N–F O–F C–F b) C–F N–O Si–F c) Cl–Cl B–Cl S–Cl
Arrange the following bonds from most to least polar: 
a) N–F O–F C–F
b) C–F N–O Si–F
c) Cl–Cl B–Cl S–Cl

9. Dipole Moment
Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
Use an arrow to represent a dipole moment.
Point to the negative charge center with the tail of the arrow indicating the positive center of charge.

10. Water - polarity, dipole, and electrostatic potential map
On electrostatic potential map red indicates electron-rich regions and blue indicates electron-poor regions.

11. Molecules with polar covalent bonds, but no dipole moment
These geometries are symmetric in 3-dimensions

12. Ionic size - (revisited)
Negative ions gain e- additional e- repulsion, so ion > atom
Positive ions lose e- less e- replusion & many times becomes lower n shell, so ion < atom
Isoelectronic ions - ions with same number of electrons, ie. O2-, Ne, Na1+
Size?

13. Ionic bonding
Positive ions attracted to negative ions - How strongly are they bound? A total of all the electrostatic interactions in the crystalline structure.

14. Lattice Energy
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
k = proportionality constant
Q1 and Q2 = charges on the ions
r = shortest distance between the centers of the cations and anions

15. LiF lattice energy diagram
Note that formation of LiF from its elements is largely exothermic (blue shaded region).
Also, the energy released when an e- is added to F is NOT enough to remove an e- from Li. The lattice energy (attraction of ions) is dominant force in forming ionic compounds.

16. MgO v. NaF lattice energy diagram
Note that the electron affinity of O -> O-2 is endothermic! Adding 1st e- is exothermic, but the 2nd e- not. (blue)
Also removing 2nd e- from Mg is large endo (yellow)
Lattice E is huge 2+ & 2- ions (purple)
Overall Energy of these ions is about the same!!!
A Variety of factors determine composition of ionic compounds
Why not Na+2 and F-2 ?

17. Calculation for % Ionic Character
None of the bonds reaches 100% ionic character even with compounds that have a large electronegativity difference. A nice model and calculation, but these are all diatomic. What to do about polyatomics NH4NO3 etc.?
What are we going to do?
Practical solution… any compound that conducts electricity when melted is classified as ionic.

18. Fundamental Properties of Models
Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
A model does not equal reality.
Models are oversimplifications, and are therefore often wrong.
Models become more complicated and are modified as they age.
We must understand the underlying assumptions in a model so that we don’t misuse it.
When a model is wrong, we often learn much more than when it is right.

19. Covalent Bonding Recall
A bond is a combination of forces that hold groups of atoms together and make them function as a unit.
A bond will form if the energy of the aggregate is lower than that of the separated atoms.

20. Covalent bond energies
To break bonds, energy must be added to the system (endothermic).
To form bonds, energy is released (exothermic).
ΔH = Σn×D(bonds broken) – Σn×D(bonds formed)
D represents the bond energy per mole of bonds (always has a positive sign).

21. Bond Energy example Predict ΔH for the following reaction:
Given the following information:
Bond Energy (kJ/mol)
C–H
C–N
C–C
891
ΔH = –42 kJ

22. Bond Energy Problem #66 pg395
Using Table 8.4 Bond energies estimate ΔH for the following reaction
CH3OH (g) + CtO (g) --> CH3COOH(l)
bonds broken - bonds formed
(3C-H, C-O, O-H, CtO) - (3C-H,C-C, C=O, C-O, O-H )
ΔH = -20 kJ

23. Localized Electron Bonding Model
A molecule is atoms bound together by sharing e- pairs using atomic orbitals of bound atoms.
LE pairs are either lone pairs (on atoms) or bonding pairs (space between atoms)
Valence e- arrangement using Lewis structures
Prediction of geometry using VSEPR
Description of type of atomic orbitals used to share e- or hold lone pairs (details in chpt 9)

24. Lewis Structures
Procedure based on observations of 1000’s of molecules
Most critical feature is atoms achieve noble gas configuration of electrons
Use valence e- only
Total the valence e- of all atoms - do not worry about where the e- came from
Form a single bond between each pair of atoms
Arrange remaining e- to satisfy duet or octet rule

25. Lewis Structure practice
Draw Lewis structure for the following:
HF CH4
N2 NO+
CF4 NH3

26. Exceptions to Octet rule
Odd number of e- (more on this later)
ex. NO
Boron (pairs each of its 3 --> 6 valence)
ex. BF usually reacts with lone pairs
Phosphorus and Sulfur (expands octet using empty valence d-orbitals)
P pairs 5e- giving 10 valence e-
S pairs 6e- giving 12 valence e-

27. Octet exception examples
When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, place the extra electrons on the central atom.
SF4 = 34e– AsBr5 = 40e–

28. Exception rules
C, N, O, and F should always be assumed to obey the octet rule.
B and Be often have fewer than 8 electrons around them in their compounds.
Second-row elements never exceed the octet rule.
Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals.
When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond).

29. Resonance Structures
More than one valid Lewis structure can be written for a particular molecule.
NO3– = 24e–
Actual structure is an average of these 3, so bond length is not single of double, but 3 equal 1.33 bonds.

30. Resonance Structures
NO3-1 SO3
NO NO2 I3-1

31. Formal Charge Used to evaluate nonequivalent Lewis structures.
Atoms in molecules try to achieve formal charges as close to zero as possible.
Any negative formal charges are expected to reside on the most electronegative atoms.

32. Formal Charge (cont.)
Formal charge = (# valence e– on free atom) – (# valence e– assigned to the atom in the molecule).
Assume:
Lone pair electrons belong entirely to the atom in question.
Shared electrons are divided equally between the two sharing atoms.

33. Formal Charge Rules To calculate the formal charge on an atom:
Take the sum of the lone pair electrons and one-half the shared electrons.
Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom.
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species.

34. Formal charge example P: 5 – 4 = +1 O: 6 – 7 = –1 Cl: 7 – 7 = 0
Consider the Lewis structure for POCl3. Assign the formal charge for each atom in the molecule. 
P: 5 – 4 = +1
O: 6 – 7 = –1
Cl: 7 – 7 = 0
BF3 revisted - why not a double bond of one of the fluorines to boron to satisfy octet? It would also have resonance structure.
Formal charge issues!!! Lets do it.

35. VSEPR guidelines Draw the Lewis structure for the molecule.
Count the electron pairs and arrange them in the way that minimizes repulsion (put the pairs as far apart as possible.
Determine the positions of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms).
Determine the name of the molecular structure from positions of the atoms.

36. VSEPR bond angles
Lone pair takes up more space than bonding pair, so “normal” 109.5o is slightly reduced.

37. 4 pairs and 5 pairs Geometries

38. E- pair minimum repulsion shapes

39. Ball & Stick Shapes

40. Trigonal bipyramid and Octahedral angles

41. E- pair arrangements

42. Bond Length (pm) Energy (kJ/mol) Data
Bond Length Energy Bond Length Energy
H--H H--C
C--C H--N
N--N H--O
O--O H--F
F--F H--Cl
Cl-Cl H--Br
Br-Br H--I
I--I
C--C
C--C C=C
C--N CºC
C--O
C--F O--O
C--S O=O
C--Cl
C--Br N--N
C--I NºN

43. Bond Length Data (cont)
Bondlengths are determined by X-ray diffraction of solids, by electron diffraction, and by spectroscopic methods (study the light absorbed or emitted by molecules).
The bondlengths ranges from the shortest of 74 pm for H-H to some 200 pm for large atoms, and the bond energies depends on bond order and lengths. sHalf of the bondlength of a single bond of two similar atoms is called covalent radius. The sum of two covalent radii of two atoms is usually the single bondlength. For example, the covalent radii of H and C are 37 and 77 pm respectively. The C-H bond is thus (37+77) 114 pm. Note that 77 pm = 154/2 pm.