1. Example: CO3 , carbonate ion
RESONANCE STRUCTURES
For certain molecules or molecular ions, two or more EQUIVALENT Lewis structures can be drawn.
The equivalent structures have the same formal charge and are equal in energy. 2-
Example: CO3 , carbonate ion C O 1 2 3 -2

2. These equivalent structures are called RESONANCE STRUCTURES.
Three equivalent structures for the carbonate ion – all have the same formal charge. C O 1 2 3 -1
These equivalent structures are called RESONANCE STRUCTURES.
The double bond electron pair is DELOCALIZED over all three C-O bonds and hence all three C-O bonds are between pure-single and pure-double bonds.
Note: This differs from the two structures for N3- which had different bonding arrangements for the atoms, and hence not equivalent.

3. The carbonate ion is NOT a mixture of the three structures, nor does any one of the three exist independently.
The real structure is an average of the three structures, and properties of these bonds are such that they are intermediate between single and double bonds.
All three C-O bonds identical, the bond lengths are between a single bond between C and O and a double bond between C and O
Other examples of molecules that have more than one equivalent resonance structures are NO3-, SO3

4. EXCEPTIONS TO THE OCTET RULE
Although the octet rule is very useful, it does not apply in all cases. There are three classes of exceptions
1) Compounds with odd number of electrons
2) Compounds with a shortage of electrons, less
than the octet.
3) Compounds with an abundance of electrons,
exceeding the octet

5. 1) Molecules with ODD number of valence electrons
An example is nitrogen monoxide, NO.
Molecules with an odd number of electrons cannot satisfy the octet rule for all its atoms.
In this case the octet rule must be given up for one of the atoms by leaving an unpaired lone electron.
ALL BONDING ELECTRONS MUST BE PAIRED.

6. NO is stable molecule with 11 valence electrons
N=O N=O
(i)
(ii)
Two Lewis structures can be drawn for NO
In (i) N has an octet, but not O; in (ii) O has an octet but not N
Calculate the formal charge of each atom for (i) and (ii) to determine which is the more favored structure.
N=O N=O
(i)
(ii) -1 +1
Structure (ii) is favored since the formal charges on both atoms is zero; in (i) N is -1 and O is +1, with the more electronegative atom, O, possessing a +1 formal charge

7. 2) Electron Deficient Molecules Example BF3
Lewis model predicts the following structure for BF3 B F +1 -1
However, experimental evidence suggests that the structure of BF3 is: B F

8. B F +1 -1
(i)
(ii)
(i) assigns a +1 formal charge to a very electronegative atom which is undesirable
(ii) is the observed structure, even though B does not have an octet, the formal charges on all atoms are zero.

9. 3) Valence-Shell Expansion
The elements in Groups IIIA to VII A in the third period and beyond show a tendency to surround themselves with more than 8 electrons.
It is possible for these elements through “valence shell expansion” for the central atom to have more than 8 electrons
In cases like this, after accounting for bonding electrons, assign lone pairs to the outer atoms to give them octets.
If any electrons still remain, assign them to the central atoms as lone pairs.

11. Example: SF6
Number of valence electrons = 6 + (6x7) = 48
Total number of electrons needed for each atom to satisfy an octet = 7x8 = 56
Number of shared electrons = = 8
8 electrons, 4 bonds, is not enough bonds for SF6
Hence, assign one bond pair to each bond, which for SF6 is 6.
Then assign lone pairs to the outer atoms to give them an octet. If any electrons remain, assign them to the central atom. S F

12. Count all electrons (6x6) + (6x2) = 48 electronsF
Count all electrons
(6x6) + (6x2) = 48 electrons
Formal charge on all atoms: 0

13. POCl3 - phosphoryl chloride-1 +1 P Cl O
(ii)
(i)
i) non zero formal charges
ii) valence shell expansion for P

14. Shapes of Molecules
Molecules have definite shape i.e. definite geometries.
For example, H2O is bent and CO2 is linear. O H
104.5 o C =
180o

15. Lewis stuctures do not provide information on a molecule’s geometry, i
Lewis stuctures do not provide information on a molecule’s geometry, i.e. shape.
VALENCE SHELL ELECTRON-PAIR REPULSION (VSEPR) THEORY: predicts molecular shapes based on Lewis stuctures
VSEPR derives from the realization that bonding and non-bonding electron pairs will locate themselves in space so as to minimize repulsive interactions.
The geometry of the electron pairs, and the covalent bonds that form, can be predicted using VSEPR

16. Molecular geometry depends on the total number of electron pairs
To use the VSEPR theory to predict molecular geometry, need to know how electron pairs can be arranged around a central atom so as to minimize electron-electron repulsion.
An electron pairs arranges itself in three-dimensional space such that it is equally far apart from all the other electron pairs.
The resulting arrangement of the electron pairs in a molecule depends on the NUMBER of electron pairs.
Note: For VSEPR, the electron pairs we refer to are the valence electrons

17. For a central atom surrounded by TWO electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs is LINEAR
180o
LINEAR
For a central atom surrounded by THREE electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs is TRIGONAL PLANAR
120o
TRIGONAL PLANAR

18. For a central atom surrounded by FOUR electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs is TETRAHEDRAL
109.47o
TETRAHEDRAL
For a central atom surrounded by FIVE electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs TRIGONAL BIPYRAMIDAL
TRIGONAL BIPYRAMIDAL
90o
120o

19. For a central atom surrounded by SIX electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs OCTAHEDRAL
90o
OCTAHEDRAL
We have been talking about the location of the electron pairs about the central atom.
This maybe different from molecular geometries i.e. the structure or shape of a molecule

20. If O is bonded to 2 H atoms, why is the structure of water not linear?
Water has a “BENT” structure with the H-O-H angle measured to be 104.5o. O H
104.5 o
If O is bonded to 2 H atoms, why is the structure of water not linear?

21. MOLECULAR GEOMETRY
Molecular geometry is the arrangement of ATOMS about the central atom.
If there are no non-bonding electrons, then the electron geometry and the molecular geometry is the same.
If there are non-bonding electrons on the central atom then the molecular geometry differs from the electron geometry but is obtained from knowledge of the electron geometry about the central atom.

22. To determine the molecular geometry about the central atom of a molecule:
1) Write the best Lewis structure for the molecule or ion
2) Count the total number of ATOMS attached to the central atom plus the number of NONBONDING ELECTRON PAIRS on the central atom .
3) Using the VSEPR geometries establish the ELECTRON geometry about the central atom.
4) Finally examine the placement of atoms and identify the molecular geometry.

23. Step 1: Write the best Lewis structure for the molecule
Total number of valence electrons = = 8
Number of electrons needed for each to complete an octet (remember for H need two electrons) = = 12
Shared electrons = 12- 8=4 => two bonds
Number of unpaired electrons = 4 => 2 lone pairs on O
Lewis structure of H2O is: H O

24. 2) Count the total number of atoms attached to the central atom plus the number of non-bonding electron pairs on the central atom H O
For water,
Total number of electron pairs = pairs of bonding electrons (bp) + 2 pairs of non-bonding electrons (np) = 4 pairs

25. 3) Then using the VSEPR geometries establish the electron geometry about the central atom.
According to VSEPR theory: For a central atom surrounded by four electron pairs, the optimum geometry which minimizes repulsion between the electrons pairs is tetrahedral
Arrange the two nonbonding electron pairs and the two H atoms at the apices of the tetrahedron, with O at the center O H

26. 4) Finally examine the placement of atoms and identify the molecular geometry.
To determine the molecular geometry of water, look at the arrangement of just the atoms (in this case, H,O,H) O H
Looking at the arrangement of the atoms in H2O, we see that the molecule is bent.

27. For a tetrahedral geometry, expect that the H-O-H angle is 109.5o
The measured H-O-H angle is 104.5o, slightly smaller than predicted.
Non-bonding electron pairs are more diffuse than bonding electron pairs, taking up more space.
Model

28. Draw the Lewis structure to determine the number of bonding electron pairs and number of nonbonding pairs on the central atom (a double or triple bond counts as a single electron pair when using VSEPR theory).
Arrange the bonding and nonbonding electron pairs around the central atom in a geometry predicted by VSEPR.
Position the atoms bonded to the central atom where the bonding electrons are positioned.
Determine the molecular geometry based on the arrangement of the atoms

29. Two electron pairs Example: CO2
Lewis structure:
O=C=O
Number of electron pairs around C = 2 (count the double bond as a single electron pair).
Both are bonding pairs
VSEPR predicts a linear molecular geometry for CO2
Model

30. Three electron pairs B F BF3 – Lewis structure
BF3 – Lewis structure
The three electron pairs around the central B atom are all bonding pairs.
According to VSEPR the three electron pairs around the central atom are positioned at the vertices of an equilateral triangle
In the case of BF3, since the three electron pairs are bonding pairs, the three F atoms are positioned at the vertices of the equilateral triangle.
Model

31. F B trigonal planar NO2- Lewis structure: O=N-O
For VSEPR count two bonding electron pairs and one nonbonding electron pair.
Hence the three electron pairs are situated at the vertices of an equilateral triangle.

32. The molecular geometry if NO2- is not trigonal planar, but bent.
The two O atoms are situated at two of the vertices and the nonbonding electron pair at the third.
The molecular geometry if NO2- is not trigonal planar, but bent.
The O-N-O angle is less than 120o due to the presence of the nonbonding pair on N N O
>120o
bent -1
<120o

33. methane, CH4 tetrahedral water, H2O bent ammonia, NH3 pyramidal
Four electron pairs
For molecules with four electron pairs, the electron pairs surround the central atom in a TETRAHEDRAL geometry.
The molecular geometry depends on the number of bonding and nonbonding electron pairs on the central atom.
Examples: C H
methane, CH4
tetrahedral O
water, H2O
bent N
ammonia, NH3
pyramidal
Model

34. Electronic geometry is trigonal bipyramidal.
Five Electron Pairs
Electronic geometry is trigonal bipyramidal.
However, unlike the 2, 3 and 4 electron geometries all vertices of a trigonal bipyramid are NOT equivalent.
There are two AXIAL positions and three EQUATORIAL positions.
90o
120o
axial
equatorial

35. axial-central atom-axial = 180o axial-central atom-equatorial = 90o
Angle between
axial-central atom-axial = 180o
axial-central atom-equatorial = 90o
equatorial -central atom- equatorial = 120o
The positions of the nonbonding electron pairs relative to each other and the bonding pairs have to account for the non-equivalence of the axial and equatorial positions.
To position the nonbonding and bonding electrons: nb-nb repulsion > nb-b > b-b

36. Hence, position nonbonding pairs first, then bonding pairs accounting for the repulsion between the electron pairs.
PCl5 P Cl Cl P
trigonal
bipyramidal
Model

37. SF4 S F
Position the nonbonding electron pair on S at the equatorial position since this minimizes nb-b repulsions F F S F
“see-saw” F

38. Angles in SF4 are different from the predicted 180o, 120o and 90o due to the non-bonding electron pair F S
187o
87o
102o
Other examples
ClF3 - distorted T
XeF2 - linear
Model

39. Six electron pairs Octahedral geometry
All vertices are the equivalent, hence bonding and nonbonding electron pairs can be positioned at any vertex
OCTAHEDRAL
Molecular geometry determined by number and position at atoms and nonbonding electron pairs
Model

40. Examples S F
SF6
octahedral Br F
BrF5
square pyramidal Xe F
XeF4
square planar

41. Molecular Properties
Bond Lengths
The bond distance is the distance between the two nuclei of the two atoms bonded together.
Bond lengths depend on the elements and whether the bond between the two elements is single, double or triple
Typically:
single bond length > double bond length > triple bond length

42. Bond Energies
If two atoms in a diatomic molecule are pulled far enough apart the bond between the atoms breaks.
The energy needed to break the bond is called the bond energy or bond dissociation energy.
Bond energies depend on the nature of the atoms bonded together and whether the bond is single, double or triple.
Typically single bond energies < double bond energy < triple bond energy

43. Dipole Moments
The bonding electron pair in HCl is not shared equally between the H and Cl atoms.
This is because Cl is more electronegative than H, and tends to attract the electron pair closer towards it.
Hence the electron distribution for the bonding pair of electrons is not evenly distributed over both the H and Cl atoms, but is more concentrated over the Cl atom.
The more electronegative atom develops a slight negative charge and the more electropositive atom a slight positive charge

44. This charge separation results in a DIPOLE MOMENT, and a POLAR HCl bond.
The dipole moment is defined as the product of the magnitude of the separated charges and the distance between the centers of the charges.
The dipole moment is a vector H Cl

45. The larger the electronegative difference between the bonded atoms, larger is the magnitude of the dipole moment and hence more polar is the bond.
All heteronuclear diatomic molecules (A-B. where A≠B) are polar molecules with the degree of polarity depending on the electronegativity differences between the two atoms.
Heteronuclear diatomics have non-zero dipole moments
All homonuclear diatomic molecules (A2) are non-polar, with zero dipole moments

46. Polyatomic molecules
Bonds between atoms of different elements will each have a dipole moment associated with them.
The magnitude of the dipole moment of a bond depends on the electronegativity differences between the bonded atoms.
However, whether the molecule as a whole has a dipole moment depends on the geometry of the molecule.
The dipole moment is a vector and hence the dipole moment of the molecule is the vector sum of the dipole moments of each bond in the molecule.

47. Methane, CH4 C H
The net dipole moment for CH4 is zero

48. Water, H2O O H O H
Model