1. Chemical Bonding

2. Properties depend on: Type of bonds between atoms
intramolecular
Shape of the molecules
Interactions between molecules
intermolecular

3. Bond Types Metallic (metal – metal) Ionic (metal – nonmetal)
Covalent (nonmetal- nonmetal)
Key difference between the bonds is in the nature of the “positive/negative” attraction

4. Metallic Bonding Metals tend to form cations (lose electrons)
Cations and “free electrons” are said to form a “cloud or sea of electrons” in which all atoms are able to share electrons

5. The electrons are not bound to an individual atom.

6. This atomic sharing of electrons is the “glue” that holds together the metal atoms.

7. Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

9. What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged over to B's end of the bond.
Ions have been formed.

10. Ionic Bonding
Typically formed between metals and nonmetals (High difference in electronegativities)
Electrons are transferred from one atom to another (from metal to nonmetal) resulting in the formation of positive and negative ions.
The electrostatic attractions between the positive and negative ions hold the compound together.

12. Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ
Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s) DHºf = kJ

13. Ionic Bonding
Na(s) + ½Cl2(g)  NaCl(s) DHºf = kJ

14. Calculation of enthalpy of formation
The loss of an electron from an element
Ionization energy
Na(g)Na+(g) + 1e- ∆H = 496 kj/mol
The gain of an electron by a nonmetal
Electron affinity
Cl(g) + 1e-Cl-(g) ∆H = -349 kj/mol
Attraction of cation and anion
Lattice energy
Na+(g) + Cl-(g)NaCl(s) ∆H = -788 kj/mol

15. Ionic Bonding Energetics of Ionic Bond Formation
Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions.
k is a proportionality constant
(depends on solid structure & e- configs of ions),
Q1 and Q2 are the charges on the ions
d is the distance between ions.
Charge
Distance El

16. due to charge increase
due to size increase

17. What happens if two atoms of equal electronegativity bond together?
If the atoms are equally electronegative, the bonding electrons are evenly shared

18. What happens if B is slightly more electronegative than A?
B will attract the electron pair rather more than A does.
Uneven sharing results in one side of the bond being more negative than the other (polarity)

19. Covalent Bonding Typically between two or more nonmetals
No, or low, difference between electronegativities
Positive nucleus is attracted to negative electron cloud of other atom

21. Nonpolar Covalent Both atoms have the same electronegativity.
Usually two identical atoms, for example, H2 or Cl2 molecules
This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms.

23. Polar Covalent
A covalent bond in which there is a separation of charge between one end and the other - in other words in which one end is slightly positive and the other slightly negative.

25. Covalent Bonding Lewis Structures
Covalent bonds can be represented by the Lewis symbols of the elements:
In Lewis structures, each pair of electrons in a bond is represented by a single line:

26. Covalent Bonding Multiple Bonds
It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds):
One shared pair of electrons = single bond (e.g. H2);
Two shared pairs of electrons = double bond (e.g. O2);
Three shared pairs of electrons = triple bond (e.g. N2).
Generally, bond distances decrease as we move from single through double to triple bonds.

27. Covalent Bonding Multiple Bonds
Generally, bond distances decrease as we move from single through double to triple bonds.
Generally, bond energies increase as we move from single through double to triple bonds.

28. In order to determine
# bonds, we need to learn
how to draw molecules!

29. Bond Energy and Enthalpy
The enthalpy of a reaction depends on the strength of the bonds of the molecules involved in the reaction.
-A reaction with tightly bound reactants will require a higher input of energy to make the reaction proceed than one with loosely bound reactants.
-Likewise, the amount of energy required to form the bonds of the products affects the overall enthalpy.

30. Bond Energy and Enthalpy
The enthalpy of a reaction is given by:
    
∆ H = Σbond energies of reactants -
Σbond energies of products
Note that bond energies are always positive quantities.

31. Bond Energy and Enthalpy
Example:
Using bond energies, calculate the enthalpy of the following reaction.
    C3H8(g ) + 5O2(g ) 3CO2(g ) + 4H2O(l )
H-C J/mol C-C J/mol H-O J/mol   
C=O J/mol C-O J/mol   
O=O J/mol O-O J/mol

32. Bond Energy and Enthalpy
Σ bond energies of reactants =
8(414) + 2(347) + 5(502) = 6516 kJ/mol
Σ bond energies of products =
6(730) + 8(464) = 8092 kJ/mol
The enthalpy change is therefore kJ/mol.

33. Drawing Lewis Structures
Add valence electrons.
Write symbols for the atoms
Show initial bondings.
Try to complete the octets (8e-)
If there are not enough electrons try multiple bonds.

34. Drawing Lewis Structures
NF3
Add valence electrons.
Draw skeleton structure
-put atom with lowest electronegativity
in middle (except hydrogen)
4. Show initial bondings.
5. Try to complete the octets (8e-)
6. If there are not enough electrons try multiple bonds. . N
5 + (3x7) = 26e- . F . F . F . F . F . F . N .

35. Polyatomic Ions Same procedure except: •Take charge into account
-add electrons for negative charge
-subtract electrons for positive charge
•Brackets around structure with charge shown in upper right
[ structure ]charge

36. Resonance Structures
The Lewis structure of ozone (O3)

37. Resonance Structures
However... known facts about the structure of ozone:
The bond lengths between the central oxygen and the other two oxygens are identical:

38. Resonance Structures
We would expect that if one bond was a double bond that it should be shorter than the other (single) bond
Since all the atoms are identical (oxygens) which atom is chosen for the double bond?

39. Resonance Structures
-These Lewis structures are equivalent except for the placement of the electrons
-Equivalent Lewis structures are called resonance structures, or resonance forms
-The correct way to describe ozone as a Lewis structure would be:

40. Resonance Structures
The important points to remember about resonance forms are:
The molecule is not rapidly oscillating between different discrete forms

41. Resonance Structures
There is only one form of the ozone molecule, and the bond lengths between the oxygens are intermediate between characteristic single and double bond lengths between a pair of oxygens
We draw two Lewis structures (in this case) because a single structure is insufficient to describe the real structure

42. Exceptions to Octet Rule
Less than an octet:
•”Wimpy” atoms bonding with highly electronegative atoms
-typical of B, Be, Al

43. Exceptions to Octet Rule
Greater than an octet:
(central atom must have a d sublevel)
-more than 4 atoms around central atom (PCl5)
-extra pairs of valence electrons (I3-)

44. Formal Charge
Sometimes when writing a Lewis structure you come across two different ways to write the molecule, both which look fine.
In this case, you should use formal charge to decide which structure is correct for the molecule.

45. Formal Charge
The formal charge is the difference in the number of valence electrons in the atom and the number of valence electrons in the Lewis structure.

46. Formal Charge Cf = Ev - (Eu + 1/2Ep)
The equation for the formal charge of any atom in a Lewis structure is
Cf = Ev - (Eu + 1/2Ep)

47. Formal Charge Cf is the formal charge
where
Cf is the formal charge
Ev is the number of valence electrons in the bare atom
Eu is the number of electrons in lone pairs on the atom in the Lewis structure
Ep is the number of electrons in bonded pairs on the atom in the Lewis structure

48. Formal Charge
To decide if a given structure is correct, check the formal charge on some atoms in all possible structures. In general the most likely Lewis structure has:
-all formal charges as close to zero as possible
-negative formal charges on electronegative atoms like halogens or oxygen.

49. Formal Charge
For example, consider the methanol molecule CH3OH. This can be written two different ways:
In both cases the octet rule is satisfied for all of the atoms in the structure. Which is correct?

50. Formal Charge For the leftmost structure
The carbon has four bonds, each worth 2 electrons, for a total of eight. It has no lone pairs. Thus, Cf = 4 - (0 + 1/2*8) =0
The oxygen has two bonds, each worth 2 electrons, for a total of four. It has two lone pairs. Thus, Cf = 6 - (4 + 1/2*4) =0

51. Formal Charge For the rightmost structure
The carbon has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 4 - (2 + 1/2*6) = -1
The oxygen has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 6 - (2 + 1/2*6) = +1

52. Formal Charge
The leftmost structure has the formal charges closer to zero, and thus is probably the correct structure.

53. A water molecule is angular or bent.
Molecular Geometry…
…is simply the shape of a molecule.
Molecular geometry is found using the Lewis structure, but the Lewis structure itself does NOT necessarily represent the molecule’s shape.
A water molecule is angular or bent.

54. VSEPR
Valence-Shell Electron-Pair Repulsion (VSEPR) is a simple method for determining geometry
Basis: pairs of valence electrons in bonded atoms repel one another.
These mutual repulsions push electron pairs as far from one another as possible.
B B B A
B A B B

55. Electron Geometries
An electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons.
The mutual repulsions among electron groups lead to an orientation of the groups that are called electron geometry.
These geometries are based on the number of electron groups: 2
linear 3
trigonal planar 4
tetrahedral 5
trigonal bipyramidal 6
octahedral

56. A Balloon Analogy Each electron group may be:
-an unshared pair of electrons, or
-a bond (single, double, triple bonds are each counted as one electron group).

57. VSEPR Notation
In the VSEPR notation used to describe molecular geometries, the central atom in a structure is denoted as A, terminal atoms as X, and the lone pairs of electrons as E.
Example: ClF3 is designated AX3E2. It has three groups (atoms in this case) around the Cl atom, and two lone pairs of electrons on the Cl (draw the Lewis structure to see).

62. Polar Molecules And Dipole Moments
A polar bond is a bond with separate centers of positive and negative charge.
A molecule with separate centers of positive and negative charge is a polar molecule.
The dipole moment (m) of a molecule is the product of the magnitude of the charge (d) and the distance (d) that separates the centers of positive and negative charge.
m = dd
A unit of dipole moment is the debye (D).
One debye (D) is equal to 3.34 x C m.

63. Polar Molecules In An Electric Field
An electric field causes polar molecules to “line up” but has no effect on nonpolar molecules.

64. Bond Dipoles And Molecular Dipoles
A polar covalent bond has a bond dipole; a separation of positive and negative charge centers in an individual bond.
Bond dipoles have both a magnitude and a direction (they are vector quantities).
A molecule can have polar bonds, but may be a nonpolar molecule – IF the bond dipoles cancel.

65. Bond Dipoles And Molecular Dipoles
CO2 has polar bonds, but is a linear molecule; the bond dipoles cancel and it has no net dipole moment (m = 0 D)
No net dipole
The water molecule has polar bonds also, but is an angular molecule.
The bond dipoles do not cancel, so water is a polar molecule.
Net dipole

66. Molecular Shapes And Dipole Moments
Molecular polarity can be predicted based on the following three-step approach:
Use electronegativity values to predict bond dipoles.
Use the VSEPR method to predict the molecular shape.
From the molecular shape, determine whether bond dipoles cancel to give a nonpolar molecule or combine to produce a resultant dipole moment for the molecule.

67. DETERMINING MOLECULAR POLARITY
1. Molecules are non-polar or polar
2. Non-polar molecules have an even (symmetrical) distribution of charge (+ or – )
If all atoms are the same in a 2-atom molecule (non-polar bonds)
(H2 , N2 , Br2 )
If there are no lone pairs on the central atom and the attached atoms are all the same
(CO2 , BCl3 , CH4)
3. Polar molecules have an uneven (assymetrical) distribution of charge. The molecule has a dipole (+ side and a – side, like a bar magnet)
If the outer atoms are different from each other
(HCl , H2CO , CH3F) OR
If there are lone pairs on the central atom
(H2O , NH3 , SO2 , O3)

68. Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms
Molecular Structure (Lewis Structures/VSEPR)
Bonds within the molecule
Valence Bond Theory
Hybridization of orbitals

69. Atomic Orbital Overlap
Valence Bond (VB) theory states that a covalent bond is formed when atomic orbitals (AOs) overlap.
In the overlap region, electrons with opposing spins produce a high electron charge density.
Overlap region between nuclei has high electron density
In general, the more extensive the overlap between two orbitals, the stronger is the bond between two atoms.

70. The measured bond angle in H2S is 92°; good agreement
Bonding In H2S
The measured bond angle in H2S is 92°; good agreement
Hydrogen atoms’ s-orbitals can overlap with the two half-filled p- orbitals on sulfur.

71. Points of VB Theory
Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms.
Bonding electrons are localized in the region of AO overlap.
For AOs with directional lobes (such as p-orbitals), maximum overlap occurs when the AOs overlap end to end.
VB theory is not without its problems…

72. Hybridization Of Atomic Orbitals
VB theory: carbon should have only two bonds, and they should be about 90° apart.
Reality: carbon has four bonds, which (singly bonded) are about 109° apart.
We can hybridize the four orbitals; mathematically combine the wave functions for the 2s orbital and the three 2p orbitals on carbon.
The four AOs combine to form four new hybrid AOs.
The four hybrid AOs are equivalent, and each has a single electron (Hund’s rule).
Four equivalent hybrid orbitals
can now form four bonds

73. sp3 Hybridization
Hybridizing an s-orbital with three p-orbitals gives rise to four hybrid orbitals called sp3 orbitals.
The number of hybrid orbitals is equal to the number of atomic orbitals combined.
The four hybrid orbitals, being equivalent, are about 109° apart.

74. The sp3 Hybridization Scheme
Four AOs…
…form four new hybrid AOs.

75. Methane and Ammonia Four sp3 hybrid orbitals: tetrahedral
Four electron groups: tetrahedral
Coincidence? Hardly…
In ammonia, one of the hybrid orbitals (top) contains the lone pair that is on the nitrogen atom
In methane, each hybrid orbital is a bonding orbital

76. sp2 Hybridization
Three sp2 hybrid orbitals are formed from an s-orbital and two p-orbitals.
The empty p-orbital remains unhybridized. It may be used in a multiple bond.
The sp2 hybrid orbitals are in a plane, 120o apart.
This distribution gives a trigonal planar molecular geometry, as predicted by VSEPR.

77. The sp2 Hybridization Scheme in Boron
A 2p orbital remains unhybridized.
Three AOs combine to form…
…three hybrid AOs

78. sp Hybridization
Two sp hybrid orbitals are formed from an s-orbital and a p-orbital.
Two empty p-orbitals remains unhybridized; the p-orbitals may be used in a multiple bond.
The sp hybrid orbitals are 180o apart.
The geometry around the hybridized atom is linear, as predicted by VSEPR.

79. sp Hybridization in Be
Two unused p-orbitals

80. Hybrid Orbitals Involving d Subshells
This hybridization allows for expanded valence shell compounds.
By hybridizing one s, three p, and one d-orbital, we get five sp3d hybrid orbitals.
This hybridization scheme gives trigonal bipyramidal electron-group geometry.
By hybridizing one s, three p, and two d-orbitals, we get five sp3d2 hybrid orbitals.
This hybridization scheme gives octahedral geometry.

81. The sp3d and sp3d 2 Hybrid Orbitals

82. Predicting Hybridization Schemes
In the absence of experimental evidence, probable hybridization schemes can be predicted:
Write a plausible Lewis structure for the molecule or ion.
Use the VSEPR method to predict the electron-group geometry of the central atom.
Select the hybridization scheme that corresponds to the VSEPR prediction.
Describe the orbital overlap and molecular geometry.

83. Hybrid Orbitals and Their Geometric Orientations

84. Hybrid Orbitals And Multiple Covalent Bonds
Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma (s) bonds.
All single bonds are sigma bonds.
A bond formed by parallel, or side-by-side, orbital overlap is called a pi (p) bond.
A double bond is made up of one sigma bond and one pi bond.
A triple bond is made up of one sigma bond and two pi bonds.

85. VB Theory for Ethylene, C2H4
σ-bond: overlap of s-orbital of hydrogen and sp2 hybrid orbital.
π-bond has two lobes (above and below plane), but is one bond. Side overlap of 2p–2p.
σ-bond: sp2 - sp2 overlap

86. Summary of VB theory of Ethylene

87. VB Theory: Acetylene σ-bond: s - sp overlap σ-bond: sp - sp overlap
Two π-bonds (above and below, and front and back) from 2p–2p overlap…
σ-bond: sp - sp overlap
…form a cylinder of π-electron density around the two carbon atoms.