1. 3. Count the total number of valence electrons in the atoms involved.
The total number of valence electrons is equal to the sum of the valence electrons of all the atoms.
For cations, the number of electrons is equal to sum of the valence electrons of all the atoms minus the number of electrons equal to the positive change.

2. 3. Count the total number of valence electrons in the atoms involved.
The total number of valence electrons is equal to the sum of the valence electrons of all the atoms.
For cations, the number of electrons is equal to sum of the valence electrons of all the atoms minus the number of electrons equal to the positive change.
For anions, the number of electrons is equal to sum of the valence electrons of all the atoms plus the number of electrons equal to the negative change.

3. 4. In molecules where there is a central atom bonded to two or more atoms, first draw a single bond between the central atom and each of the surrounding atoms in such a way that the octet rule is met for all the surrounding atoms.

4. 4. In molecules where there is a central atom bonded to two or more atoms, first draw a single bond between the central atom and each of the surrounding atoms in such a way that the octet rule is met for all the surrounding atoms. If the octet rule is not met for the central atom then use the lone pairs on the surrounding atoms to form double and triple bonds with the central atom.

5. 4. In molecules where there is a central atom bonded to two or more atoms, first draw a single bond between the central atom and each of the surrounding atoms in such a way that the octet rule is met for all the surrounding atoms. If the octet rule is not met for the central atom then use the lone pairs on the surrounding atoms to form double and triple bonds with the central atom. As a final check, make sure the total number of electrons in chemical bonds and lone pairs is equal to the total number of valence electrons of the atoms.

9. Exceptions to the Octet Rule

10. Exceptions to the Octet Rule
The octet rule helped to advance our knowledge of chemical bonding in the 1920’s and 1930’s.

11. Exceptions to the Octet Rule
The octet rule helped to advance our knowledge of chemical bonding in the 1920’s and 1930’s.
It was soon realized that the octet rule had its limitations and could not account for the bonding in several types of compounds.

12. 1. Compounds with more than 8 electrons on the central atom.

13. 1. Compounds with more than 8 electrons on the central atom.
Cases where the central atom is bonded to more than four atoms will fall into this group.

14. 1. Compounds with more than 8 electrons on the central atom.
Cases where the central atom is bonded to more than four atoms will fall into this group.
Example: SF6 The Lewis structure is

15. 1. Compounds with more than 8 electrons on the central atom.
Cases where the central atom is bonded to more than four atoms will fall into this group.
Example: SF6 The Lewis structure is S

16. In SF6 there are a total of 12 electrons around the central S atom, so this is an exception to the octet rule.

17. In SF6 there are a total of 12 electrons around the central S atom, so this is an exception to the octet rule.
If there are more than 4 four bonds from the central atom, the central atom will not satisfy the octet rule.

18. In SF6 there are a total of 12 electrons around the central S atom, so this is an exception to the octet rule.
If there are more than 4 four bonds from the central atom, the central atom will not satisfy the octet rule.
Second example: PF5

20. 2. Electron-deficient compounds
2. Electron-deficient compounds. Some compounds are said to be electron-deficient because there are fewer than 8 electrons about each atom.

21. 2. Electron-deficient compounds
2. Electron-deficient compounds. Some compounds are said to be electron-deficient because there are fewer than 8 electrons about each atom. Boron and aluminum form a number of compounds which belong to this group.

22. 2. Electron-deficient compounds
2. Electron-deficient compounds. Some compounds are said to be electron-deficient because there are fewer than 8 electrons about each atom. Boron and aluminum form a number of compounds which belong to this group.
Example: BF3 The Lewis structure is

23. 2. Electron-deficient compounds
2. Electron-deficient compounds. Some compounds are said to be electron-deficient because there are fewer than 8 electrons about each atom. Boron and aluminum form a number of compounds which belong to this group.
Example: BF3 The Lewis structure is

24. Key fact: F does not form double bonds, that is, the surrounding F atoms in BF3 do not share lone pair electrons with the central B atom.

25. Key fact: F does not form double bonds, that is, the surrounding F atoms in BF3 do not share lone pair electrons with the central B atom.
So in this case the central B atom does not satisfy the octet rule – it is only surrounded by 6 electrons. Hence, it is electron-deficient.

26. Electron-deficient compounds will generally be fairly reactive.

27. Electron-deficient compounds will generally be fairly reactive.
Example:
F H F H
F B : N H F B N H

28. Electron-deficient compounds will generally be fairly reactive.
Example:
F H F H
F B : N H F B N H
Since the boron atom is electron-deficient, it readily forms a bond with the nitrogen in ammonia.

29. Coordinate covalent bond (also called a dative bond): The pair of electrons is supplied by one of the two atoms which are bonded.

30. Coordinate covalent bond (also called a dative bond): The pair of electrons is supplied by one of the two atoms which are bonded. In the reaction of BF3 with NH3 the B N bond is a coordinate covalent bond, since the N atom provides both electrons.

31. 3. Odd-electron molecules: A small percentage of compounds contain an odd number of electrons.

32. 3. Odd-electron molecules: A small percentage of compounds contain an odd number of electrons. Such compounds will not satisfy the octet rule for each atom in the compound.

33. 3. Odd-electron molecules: A small percentage of compounds contain an odd number of electrons. Such compounds will not satisfy the octet rule for each atom in the compound.
Examples: NO, NO2, ClO2

34. 3. Odd-electron molecules: A small percentage of compounds contain an odd number of electrons. Such compounds will not satisfy the octet rule for each atom in the compound.
Examples: NO, NO2, ClO2 or
(1) (2)

35. 3. Odd-electron molecules: A small percentage of compounds contain an odd number of electrons. Such compounds will not satisfy the octet rule for each atom in the compound.
Examples: NO, NO2, ClO2 or
(1) (2)
The N atom satisfies the octet count but the O atom does not, in the first structure, and the opposite for the second structure.

36. Resonance

37. Resonance
Resonance: The situation in which two or more Lewis structures are used to represent a particular molecule.

38. Resonance
Resonance: The situation in which two or more Lewis structures are used to represent a particular molecule. For SO3 there are three ways of drawing its Lewis structure.

39. Bond length: The distance between two nuclei of two bonded atoms in a molecule.

40. Bond length: The distance between two nuclei of two bonded atoms in a molecule. In general, triple bonds are shorter than double bonds, which in turn, are shorter than single bonds, for the same pair of atoms.

41. A single Lewis structure for SO3 is S

42. A single Lewis structure for SO3 is S For a single Lewis structure for SO3 we would expect that the two S O bonds would be longer than the S O bond.

43. A single Lewis structure for SO3 is S For a single Lewis structure for SO3 we would expect that the two S O bonds would be longer than the S O bond. However, experimental evidence indicates that all three sulfur to oxygen bonds are equal in length.

44. Therefore, the Lewis structure S is not truly representative of the molecule.

45. Therefore, the Lewis structure S is not truly representative of the molecule. One way out of this dilemma is to draw all three Lewis structures to represent the molecule.

46. S S S

47. S S S

48. S S S Each of these three structures is called a resonance form or resonance structure.

49. S S S Each of these three structures is called a resonance form or resonance structure. The symbol indicates that the structures shown are resonance forms.

50. Note: None of the resonance structures exist separately.

51. Note: None of the resonance structures exist separately.
Don’t think of SO3 as actually oscillating among these three forms.

52. The important point is that the properties of SO3 can be properly accounted for only by considering all three resonance structures together. This can be difficult to represent in a single diagram.

53. Another example, nitrate ion, NO-3 N N N

54. Another example, chlorine dioxide, ClO2
(1) (2) (3)

55. Another example, chlorine dioxide, ClO2
(1) (2) (3)
Resonance structures (1) and (2) are equivalent in an energetic sense, but structure (3) has a different energy.

56. If these three resonance structures are used to represent the ClO2 molecule, the question naturally arises as to how much weight we should attach to structure (1) and (2) in comparison to structure (3).

57. If these three resonance structures are used to represent the ClO2 molecule, the question naturally arises as to how much weight we should attach to structure (1) and (2) in comparison to structure (3). There is an experimental technique that enables us to determine how much time the unpaired electron spends on each of the three atoms, so that the relative importance of the structures (1) and (2) versus structure (3) can be estimated.

58. Benzene C6H6

59. Benzene C6H6
This is a very important example in organic chemistry.

60. Benzene C6H6
This is a very important example in organic chemistry. C C C