1. Covalent Bonding
Lewis Structures

2. Octet Rule
In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases.
H has 1 electron, but a pair of H’s shares 2 electrons to form a covalent single bond resulting in a molecule of H2

3. The Octet Rule in Covalent Bonding
How many electrons would carbon, nitrogen, oxygen, and fluorine need to share in order to achieve the same electron configuration as neon?

4. Bonding for Second Row Elements
Draw electron dot structures for Li, Be, B, C, N, O and F.
Predict how many bonds each atom must form to attain a noble-gas configuration.
Can Li form a covalent bond and reach stability?
Which elements can reach stability by forming covalent bonds?
Can fluorine form an ionic bond?
Are the bonds in nitrogen molecules, N2, ionic or covalent?

5. Single Covalent Bonds
When a single pair of electrons is shared, a single covalent bond forms.
The shared electron pair, often referred to as the bonding pair, is represented by either a pair of dots or a line in the structural formula.
H:H or H-H (structural formulas)
H2 (molecular formula)

6. Pairs of electrons… Shared – between atoms (bonding pairs)
Unshared – not between atoms (lone pairs)

7. Lewis Structures
Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

8. Writing Lewis Structures
PCl3
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
(7) = 26

9. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
26  6 = 20

10. Writing Lewis Structures
Fill the octets of the outer atoms.
RECALL
To be stable, an atom in a molecule needs 8 total valence electrons
These electrons come in pairs.
The pairs can be shared with another atom, or unshared.
H and He have their own rule – the Duet Rule
Keep track of the electrons:
26  6 = 20  18 = 2

11. Writing Lewis Structures
Fill the octet of the central atom.
These are unshared pairs of electrons or lone pairs.
Keep track of the electrons:
26  6 = 20  18 = 2  2 = 0

12. Practice Lewis Electron Dot Structures
PH3 H2S HCl CCl4 SiH4

13. Single Covalent Bonds F F F F F - F F2
Group 7A will form single covalent bonds
Each Fluorine has 7 valence electrons. Each fluorine will share 1 valence electron with the other fluorine, so that each fluorine possesses a stable octet of valence electrons.
F F
F F
F - F F2

14. Single Covalent Bonds H O H H - O - H H N H H H - N - H - H H - C - H
Group 6A will form two covalent bonds.
Oxygen in water will form one single covalent bond with each
Hydrogen atom.
Group 5A will form three covalent bonds.
Nitrogen in ammonia will form one single covalent bond with
each Hydrogen atom.
Group 4A (C and Si) will form four covalent bonds.
Carbon in methane will form one single covalent bond with each
H O H
H - O - H
H N H H
H - N - H - H
H - C - H H -
H C H H

15. Representing Molecules
Practice drawing structural formulas, electron-dot structures, and orbital diagrams for the following molecules: Cl2, OF2, SCl2, and NCl3

16. Versatile Carbon
How many covalent bonds do you think carbon forms based on its electron configuration, 1s22s22p2?

17. Electron Promotion
Elements in groups 3A and 4A promote s electrons to p orbitals, increasing their bonding capacity.
Promotion of one 2s electron to the 2p orbital allows carbon to make 4 bonds rather than 2.

18. Representing Molecules
Practice drawing structural formulas, electron-dot structures, and orbital diagrams for the following molecules: CCl4, CHCl3, and C2H6

19. Multiple Bonds
When you don’t have enough valence electrons to give all atoms a NGEC (noble as electron configuration), then you might need to have a double or triple bond between atoms.
These atoms CANNOT double bond: H, F, Cl, Br, I. This is due to only one available bonding site in the shared orbital.

20. Writing Lewis Structures
6. If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.

21. Multiple Bond Elements
Carbon, Oxygen, Nitrogen, and Sulfur tend to form multiple bonds with themselves and each other.
Look for these elements, when forming LDS.
Remember that they are a “band of brothers” on the periodic table.
That is why they are CONS.

22. Practice Problems
Draw the Lewis Structure for each of the these molecules. O2 N2
HCN
CO2

23. Assigning Formal Charges
This is merely a bookkeeping method designed to determine the best LEDS. It is always best to try to obey the octet rule.
For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
Subtract that from the number of valence electrons for that atom.
The difference is its formal charge.

24. Writing Lewis Structures
The best Lewis structure…
…is the one with the fewest charges.
…puts a negative charge on the most electronegative atom.
Predict the Formal Charge for each of the following compounds:

25. Resonance Structures -
Is a structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion.
This is the Lewis structure we would draw for ozone, O3. + -

26. Resonance
But this is at odds with the true, observed structure of ozone, in which…
…both O—O bonds are the same length.
…both outer oxygens have a charge of 1/2.

27. Resonance
One Lewis structure cannot accurately depict a molecule such as ozone.
We use multiple structures, resonance structures, to describe the molecule.

28. Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.

29. Resonance Structures
More than one valid Lewis Structure can be written for a molecule or ion.
Differ only in position of electron pairs, never atoms.
Any molecule that undergoes resonance behaves as if it only has one structure.
Bond lengths are actually shorter than single bonds but longer than double bonds.

30. Resonance Structures
In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
They are not localized, but rather are delocalized.

31. Resonance Structures
The organic compound benzene, C6H6, has two resonance structures.
It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

32. Practice Problems
Draw the Lewis resonance structures for the following moleules:
SO3
SO2 O3
NO2-

33. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Ions or molecules with an odd number of valence electrons.
Ions or molecules with atoms possessing less than an octet of valence electrons.
Ions or molecules with atoms possessing more than eight valence electrons (an expanded octet).

34. Odd Number of Valence Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

35. Odd Number of Valence Electrons
A small group of molecules have an odd number of valence electrons and cannot form an octet around each atom.
NO2 and ClO2 and NO
NO2 nitrogen has 5 and each oxygen has 6 for a total of 17 electrons : . :
:O = N – O: :

36. Atom with Fewer Than Eight Valence Electrons
Consider BF3:
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.

37. Atom with Fewer Than Eight Valence Electrons
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

38. Atoms with Fewer Than Eight Valence Electrons
The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

39. Atoms with More Than Eight Valence Electrons
The only way PCl5 can exist is if phosphorus has 10 valence electrons around it.
It is allowed to expand the octet of atoms on the 3rd row or below.
Presumably d orbitals in these atoms participate in bonding.

40. Atom with More than Eight Valence Electrons
Five bonds are formed with ten electrons shared in the 1 s orbital, 3 p orbitals, and 1 d orbital of the 3rd energy level – expanded octet. P = 1s22s22p63s23p3 P 1s 2s 2p 3s 3p
The 5 red arrows are valence electrons
being shared by the 5 chlorine atoms
around the phosphorous atom. 4s 3d

41. Atom with More Than Eight Valence Electrons
Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

42. Atom with More Than Eight Valence Electrons
This eliminates the charge on the phosphorus and the charge on one of the oxygens.
The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

43. Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule by using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

44. Practice Problems
Draw the correct Lewis Dot Structures for the following compounds, which contain expanded octets:
SF6
PCl5
ClF3
[AsO4]3-

45. Coordinate Covalent Bonds
Some compounds form with fewer than eight electrons present around an atom. These are rare and highly reactive.
They will form coordinate covalent bonds with other molecules which have lone pairs of electrons to donate in order to make them stable.
H – N – H H :
H – B – H H
H – N – H H + :
H – B – H H

46. Coordinate Covalent Bond
A molecule in which the shared electron pair comes from one of the bonding atoms.
CO is an example :C:::O:
Carbon only has 4 valence electrons, so Oxygen donates one of its unshared lone pair of electrons to the carbon atom to complete its octet.

47. Practice Problems
Compare the electron dot structure for ammonia (NH3) and for a hydrogen ion (H+).
How do you explain the existence of the ammonium ion, NH4+?
What is the electron dot structure for the ammonium ion?

48. Evaluate Understanding
What is resonance?
When a molecule or ion has an odd number of valence electrons, which element carries the odd electron and why?
Why would an element have less than an octet?
What is a coordinate covalent bond?

49. Assessment
What electron configurations do atoms usually achieve by sharing electrons to form covalent bonds?
How is an electron dot structure used to represent a covalent bond?
When are two atoms likely to form a double bond between them? A triple bond?
How is a coordinate covalent bond different from other covalent bonds?
How is the strength of a covalent bond related to its bond dissociation energy?
Draw the electron resonance structures for ozone and explain how they describe bonding.
List three ways in which the octet rule can sometimes fail to be obeyed.
What kinds of information does a structural formula reveal about the compound it represents?
Draw the electron dot structures for the following molecules, which have only single covalent bonds.
H2S
PH3
ClF