2. Covalent Compounds

3. Covalent compounds contain covalent bonds Covalent bonds = sharing electrons Covalent bonds usually form between nonmetals. Covalent bonds can involve multiple pairs of electrons: single, double, triple bonds.

4. Properties of covalent compounds: Covalent compounds have low melting and boiling points. Covalent compounds are usually soft, not brittle. Covalent compounds are poor conductors.

5. Covalent Bonding Covalent bonds form by sharing electrons between atoms … … so that each atom appears to have an octet of electrons. Diatomic elements are good examples of covalent bonding.

6. The Diatomic Elements are: H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 Known as the “hairogens”: H, N & O, halogens N and O = ??? air H ogens …hence, the

7. Bonding in the Halogens F F + F  F 2 Formation of a F-F bond F

8. Bonding in the Halogens F F + F  F 2 The overlap of two p-orbitals creates the single bond between fluorine atoms. F - F F

9. See how a double bond occurs in an oxygen molecule and a triple bond occurs in a nitrogen molecule.

10. Bonding in Oxygen O O O + O  O 2 The overlap of four p-orbitals creates the double bond between oxygen atoms. O = O

11. Bonding in Nitrogen N N N + N  N 2 The overlap of six p-orbitals creates the triple bond between nitrogen atoms. N

12. Comparison of single, double and triple bonds: Bond length: s--i--n--g--l--e > d o u b l e > triple A B A B A B Bond strength: single < double < triple

13. Covalent bonds result from the overlap of orbitals.

15. Recall the shapes of the three p-orbitals z x

16. Consider two p-orbitals from two different elements: As the orbital get closer …

17. Consider two p-orbitals from two different elements: A bond occurs when the orbitals overlap end to end.

18. Gilbert N. Lewis American chemist and educator. Defined acids as electron pair acceptors and bases as electron pair donors. Explained his theory with “electron dot diagrams”. Still in use today to explain molecular structure as well as acids and bases.

19. Writing Lewis Structures 1.Add up all of the valence electrons 2.Decide on a central atom. It has the lowest EN. H is never a central atom; halogens rarely are. 3.Draw the skeleton of the molecule and connect each symbol with a dash to indicate a bonding pair of electrons

20. Writing Lewis Structures 4.Complete the octet of the terminal atoms, add all the electrons and compare to #1 5.Add any additional electrons to the central atom, even if it means having more than 8. 6.If there are not enough electrons to give every element an octet, consider multiple bonds.

21. Hydrogen can only have two electrons around it, not an octet. The central atom is frequently the one that there is only 1 of. Halogens are almost never the central atom and they never have double or triple bonds! Writing Lewis Structures Some things to remember:

22. Write the Lewis structures for the following compounds: 1. H 2 O 2. CH 4 3. OF 2 4. PCl 3 5. HCN 6. CO 7. CO 2 8. SCl 4 9. PCl 5 10. XeCl 4

23. Polyatomic Ions An ion with two or more atoms. Polyatomic ions have unique formulas and names: OH - = hydroxide ion SO 4 2- = sulfate ion PO 4 3- = phosphate ion

24. Lewis Structures of polyatomic ions Write the Lewis structure just as you would for a compound, except … …the number of valence electrons must be increased (or decreased) because of the charge on the ion. PO 4 3- 5 + 4(6) +3 = 32 electrons Consider the phosphate ion. It has three extra electrons.

25. Draw the Lewis structure of the following polyatomic ions: 1.Hydroxide ion OH - 2.Sulfate ion SO 4 2- 3.Phosphate ion PO 4 3- 4.Nitrate ionNO 3 1- 5.Ammonium ionNH 4 +