1. Covalent Bonds and Structures Chemistry 11 Ms. McGrath

2. Covalent Bonds Because covalent bonds range from nearly ionic to non-polar covalent, the shapes of molecular compounds have a great variety. We already know that ionic compounds have a crystal like shape created by the orderliness of the ions.

3. Covalent Bonds – Lewis Structures We will be analyzing and predicting the structures and properties of molecules. Lewis structures are going to be very important in determining these.

4. Covalent Bonds – Lewis Structures Draw the Lewis Structure for methanol, CH 2 O

5. Covalent Bonds – Lewis Structures Step 1 Determine the total number of valence electrons in all of the atoms in the molecule

6. Covalent Bonds – Lewis Structures Step 1 (1 C atom x 4e - ) + (1 O atom x 6e - ) + (2 H atoms x 1e - ) = 4e - + 1e - + 2e - = 7e -

7. Covalent Bonds – Lewis Structures Step 2 Draw a skeleton structure for the molecule. The atom with the largest number of unpaired electrons will be the central atom. Join the atoms with one pair of bonding electrons.

8. Covalent Bonds – Lewis Structures Step 2 O.. H:C:H

9. Covalent Bonds – Lewis Structures Step 3 Place lone pairs of electrons around all atoms, except the central atom, to form an octet of electrons. Hydrogen, of course, has only two valence electrons.

10. Covalent Bonds – Lewis Structures Step 3.. :O:.. H:C:H

11. Covalent Bonds – Lewis Structures Step 4 (a) If all the valence electrons determined in Step 1 have not been used, add one or more lone pairs around the central atom to complete an octet of electrons (b) If all of the valence electrons have been used up but the central atom does not have an octet, move one or more of the lone pairs to form double or triple bonds between the central atom and an adjacent atom

12. Covalent Bonds – Lewis Structures Step 4.. :O:.. H:C:H There are 12 electrons in the structure, which is the same as the total number in Step 1. Carbon, however, does not have an octet. Therefore, move one of the lone pairs around the oxygen atom to the position between the oxygen and carbon atoms to form a double bond.

13. Step 4 :O:.. H:C:H

14. Coordinate Covalent Bonds A coordinate covalent bond is another type of covalent bond where a pair of electrons is shared between two atoms but both electrons were originally part of one of the two atoms (i.e. both electrons of the shared pair come from one of the bonded atoms).

15. Coordinate Covalent Bonds Consider the formation of boron trifluoride (BF3):

16. Coordinate Covalent Bonds One of the boron’s orbitals is empty and the BF 3 is attracted to molecules with lone pairs of electrons that can be donated.

17. Coordinate Covalent Bonds Consider the reaction of BF 3 with NH 3 to form NH 3 BF 3 :

18. Coordinate Covalent Bonds The bond between N and B is a coordinate covalent bond – it can’t be experimentally distinguished from other shared pair covalent bonds.

19. Resonance Structures  More than one possible Lewis Structure  Draw the Lewis Structure for sulfur dioxide, SO 2

20. Resonance Structures The Lewis Structure suggests that SO 2 contains a single and a double bond. In fact, experimental measurements determine that the bond lengths in SO 2 are identical. SO 2 contains two “one and a half” bonds.

21. Resonance Structures These are models that give the same relative position of atoms as in Lewis Structures, but show different placing of their bonding pairs and lone pairs. It is important to know that resonance structures do not exist in reality.

22. Resonance Structures It is important to know that resonance structures do not exist in reality. SO 2 does not shift back and forth from one structure to another. An actual SO 2 molecule is a combination of its two resonance structure. You can imagine one pair of electrons as resonating across the entire molecule from one oxygen atom to the other.