1. Chapters
Bonding

2. Types of Bonds
Bonds tend to form in order to give each atom in the compound an Octet in its valence shell
Nonmetals with other nonmetals tend to share valence electrons, causing a covalent bond
Cl2

3. Types of Bonds
A special type of covalent bond can be seen in Network Solids, such as diamond, graphite and SiO2.
SiO2
Graphite
Diamond

4. Types of Bonds
Bonds tend to form in order to give each atom in the compound an Octet in its valence shell
Metals tend to give electrons (making themselves cations) to nonmetals (making anions), with the resulting anion attracting to the resulting cation in an ionic bond

5. Types of Bonds
Bonds tend to form in order to give each atom in the compound an Octet in its valence shell
Metals tend to release their electrons to other metals (creating metal ions attracted to a “sea of electrons”) in metallic bonding.

6. Bond Length

7. Bond Polarity
Bond polarity is caused by differences in electronegativities… the larger the difference the greater the polarity (or the greater the dipole)

8. Bond Polarity
Should the dipole become large enough, you will have an ionic bond…
Nonpolar Covalent
Polar Covalent
Ionic
Ionic

9. Effects of Polarity
If a bond is polar enough it will form not a covalent, but an ionic bond…
Nonpolar Covalent Bond
Polar Covalent Bond
Ionic Bond

10. Effects of Polarity
Polar molecules are attracted to electric fields…(bending water demo)

11. Effects of Polarity
Polar molecules are attracted to each other, and other ions…
The larger the dipole, the more ionic the character (higher melting points, etc)

12. Effects of Polarity
Rank the following compounds in order of predicted boiling points HF
HCl
HBr

13. Effects of Polarity
How does polarity affect Bond Length?

14. Effects of Polarity
“Like dissolves like” (Polar molecules will dissolve polar molecules, non-polar molecules will dissolve non-polar molecules)
More on this later

15. Effects of Polarity
So far we have looked at simple, diatomic, molecules. To look at more complex molecules, we need to observe the overall polarity of the molecule…

16. Effects of Polarity
We will learn to predict the shapes/polarities soon…

17. Formation of Binary Ionic Compounds
Lattice Energy

18. Lattice Energy for some Ionic Compounds
What causes increase in lattice energies?
Which of the following will have higher boiling points?

19. Enthalpy from bond energies
Bonds break  energy absorbed (positive)
Bonds made  energy released (negative)

20. Enthalpy from bond energies
Find the energy released or absorbed in the combustion of Methane…

22. Localized Electron Bonding Model
A model to represent where valence electrons are shared in a molecule
Electron pairs are assumed to be localized on an atom or between to atoms (bonded)… the unbonded pairs are called lone pairs
This allows us to…
Write Lewis structures
Predict molecular geometry (shape)
Describe how orbitals change (hybridize) when bonding

23. Lewis Structures
Write dots around the chemical symbol to show the number of valence electrons (only s and p electrons are shown)

24. Lewis Structures Basic Method:
Add the total number of electrons in the compound
Pair electrons to bond the atoms (avoid rings if possible)
Fill in lone pairs, and rearrange to satisfy the octet rule
Write the Lewis structure for ammonia

25. Lewis Structures Basic Method:
Add the total number of electrons in the compound
Pair electrons to bond the atoms (avoid rings if possible)
Fill in lone pairs, and rearrange to satisfy the octet rule
Write the Lewis structure for Carbon Dioxide

26. Lewis Structures Basic Method:
Add the total number of electrons in the compound
Pair electrons to bond the atoms (avoid rings if possible)
Fill in lone pairs, and rearrange to satisfy the octet rule
Write the Lewis structure for NO+

27. Lewis Structures Basic Method:
Add the total number of electrons in the compound
Pair electrons to bond the atoms (avoid rings if possible)
Fill in lone pairs, and rearrange to satisfy the octet rule
Write the Lewis structure for C2H4

28. Exceptions to the Octet Rule
Boron: can have 6 valence electrons (still prefers 8) (see boron trifluoride)
Elements in period 3 and beyond can have an exceed the octet rule (here d electrons act like valence electrons and bond)… see sulfur hexafluoride, phosporous pentachloride…

29. Lewis Structures and Resonance (and bond order)
Nitrate ion
Ozone
Benzene ring

30. Formal Charge
Formal charge is another way to “keep track” of electrons (other than oxidation state) is called formal charge. Oxidation states assume the more electronegative atom in a molecule keeps both shared electrons… formal charge assumes they are shared equally. This is more useful when evaluating if a given Lewis structure is “more valid” than another…

31. Three possible structures for CNS-
Formal Charge
Valence Electrons are assigned in an atom as such:
Lone pairs are assigned only to the local atom
Bonded pairs share electrons, thus each atom that is bonded gets one of the two bonded electrons
The sum of the formal charges must equal the overall charge on the molecule
Three possible structures for CNS-
Find the formal charge of each element above.

32. Formal Charge To determine which structure is best…
Each element should have a formal charge as close to zero as possible
The more electronegative elements will have any negative charges, if necessary
Which structure is best??

33. Three possible structures for CNO-
Formal Charge
Valence Electrons are assigned in an atom as such:
Lone pairs are assigned only to the local atom
Bonded pairs share electrons, thus each atom that is bonded gets one of the two bonded electrons
The sum of the formal charges must equal the overall charge on the molecule
To determine which structure is best…
Each element should have a formal charge as close to zero as possible
The more electronegative elements will have any negative charges, if necessary
Three possible structures for CNO-
Determine the formal charge and which structure is best

34. Molecular Shapes
The shape of a molecule plays an important role in its reactivity.
By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.
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35. What Determines the Shape of a Molecule?
Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.
Note that non bonding pairs are more repulsive than bonding pairs!!

36. Electron Domains
We can refer to the electron pairs as electron domains (aka Steric Number or Effective pairs).
In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.
The central atom in this molecule, A, has four electron domains.
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37. Valence Shell Electron Pair Repulsion Theory (VSEPR)
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”
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38. Electron-Domain Geometries
These are the electron-domain geometries for two through six electron domains around a central atom.
Note the bond angles, you will be expected to know these!!

39. Electron-Domain Geometries
All one must do is count the number of electron domains in the Lewis structure.
The geometry will be that which corresponds to the number of electron domains.
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40. Molecular Geometries
The electron-domain geometry is often not the shape of the molecule, however.
The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.
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41. Molecular Geometries
Within each electron domain, then, there might be more than one molecular geometry.
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42. Linear Electron Domain
In the linear domain, there is only one molecular geometry: linear.
NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.
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43. Trigonal Planar Electron Domain
There are two molecular geometries:
Trigonal planar, if all the electron domains are bonding,
Bent, if one of the domains is a nonbonding pair.
What are the angles between each “B” atom??

44. Nonbonding Pairs and Bond Angle
Nonbonding pairs are physically larger than bonding pairs.
Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.
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45. Multiple Bonds and Bond Angles
Double and triple bonds place greater electron density on one side of the central atom than do single bonds.
Therefore, they also affect bond angles.
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46. Tetrahedral Electron Domain
There are three molecular geometries:
Tetrahedral, if all are bonding pairs,
Trigonal pyramidal if one is a nonbonding pair,
Bent if there are two nonbonding pairs.
What are the angles between each “B” atom??
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47. Trigonal Bipyramidal Electron Domain
There are two distinct positions in this geometry:
Axial
Equatorial
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48. Trigonal Bipyramidal Electron Domain
Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.
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49. Trigonal Bipyramidal Electron Domain
There are four distinct molecular geometries in this domain:
Trigonal bipyramidal
Seesaw
T-shaped
Linear
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50. Octahedral Electron Domain
All positions are equivalent in the octahedral domain.
There are three molecular geometries:
Octahedral
Square pyramidal
Square planar
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51. Larger Molecules
In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.
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52. Larger Molecules
This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.
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53. Polarity In Chapter 8 we discussed bond dipoles.
But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.
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54. Polarity
By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.
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55. Polarity
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56. Overlap and Bonding
We think of covalent bonds forming through the sharing of electrons by adjacent atoms.
In such an approach this can only occur when orbitals on the two atoms overlap.
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57. Overlap and Bonding
Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron-electron repulsion.
However, if atoms get too close, the internuclear repulsion greatly raises the energy.
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58. Hybrid Orbitals
But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.
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59. Hybrid Orbitals Consider beryllium:
In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.
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60. Hybrid Orbitals
But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.
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61. Hybrid Orbitals
Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded as is the s orbital.
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62. Hybrid Orbitals
These two degenerate orbitals would align themselves 180 from each other.
This is consistent with the observed geometry of beryllium compounds: linear.
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63. Hybrid Orbitals
With hybrid orbitals the orbital diagram for beryllium would look like this.
The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.
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64. Hybrid Orbitals Using a similar model for boron leads to…
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65. Hybrid Orbitals …three degenerate sp2 orbitals.
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66. Hybrid Orbitals
With carbon we get…
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67. Hybrid Orbitals …four degenerate sp3 orbitals.
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68. Hybrid Orbitals
For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids.
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69. Hybrid Orbitals This leads to five degenerate sp3d orbitals…
…or six degenerate sp3d2 orbitals.
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70. Hybrid Orbitals
Once you know the electron-domain geometry, you know the hybridization state of the atom.
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71. Valence Bond Theory
Hybridization is a major player in this approach to bonding.
There are two ways orbitals can overlap to form bonds between atoms.
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72. Sigma () Bonds Sigma bonds are characterized by Head-to-head overlap.
Cylindrical symmetry of electron density about the internuclear axis.
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73. Pi () Bonds Pi bonds are characterized by Side-to-side overlap.
Electron density above and below the internuclear axis.
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74. Single Bonds
Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.
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75. Multiple Bonds
In a multiple bond one of the bonds is a  bond and the rest are  bonds.
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76. Multiple Bonds
In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps in  fashion with the corresponding orbital on the oxygen.
The unhybridized p orbitals overlap in  fashion.
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77. Multiple Bonds
In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap in  fashion to form the two  bonds.
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78. Delocalized Electrons: Resonance
When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.
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79. Delocalized Electrons: Resonance
In reality, each of the four atoms in the nitrate ion has a p orbital.
The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.
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80. Delocalized Electrons: Resonance
This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.
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81. Resonance
The organic molecule benzene has six  bonds and a p orbital on each carbon atom.
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82. Resonance
In reality the  electrons in benzene are not localized, but delocalized.
The even distribution of the electrons in benzene makes the molecule unusually stable.
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83. Molecular Orbital (MO) Theory
Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals.
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84. Molecular Orbital (MO) Theory
In MO theory, we invoke the wave nature of electrons.
If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital.
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85. Molecular Orbital (MO) Theory
If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.
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86. MO Theory
In H2 the two electrons go into the bonding molecular orbital.
The bond order is one half the difference between the number of bonding and antibonding electrons.
Bond Order represents the strength of the bond. (Big is Strong, Zero is nothing)
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87. MO Theory
For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is 1 2
(2 - 0) = 1
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88. MO Theory In the case of He2, the bond order would be 1 2 (2 - 2) = 0
Therefore, He2 does not exist.
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89. MO Theory
For atoms with both s and p orbitals, there are two types of interactions:
The s and the p orbitals that face each other overlap in  fashion.
The other two sets of p orbitals overlap in  fashion.
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90. MO Theory The resulting MO diagram looks like this.
There are both s and p bonding molecular orbitals and s* and * antibonding molecular orbitals.
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91. MO Theory
The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals.
This flips the order of the  and  molecular orbitals in these elements.
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92. Second-Row MO Diagrams
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