Covalent Bonding

1. Covalent Bonding Sec. 8.3: Molecular Structures

2. Objectives • List the basic steps used in drawing Lewis structures. • Explain why resonance occurs and identify resonance structures. • Explain 3 exceptions to the octet rule and identify molecules in which these exceptions occur.

3. Molecular Structures • Molecular structures are used to predict the arrangement of atoms in a covalent compound • 5 models are used - see pg. 253, Fig. 13 • The Lewis structure model is the most useful; letter symbols and bonds show the position of atoms & “dots” show unpaired electrons. • Electron dot diagrams must be drawn to predict the structural formula.

4. Determining Lewis Structures of Molecular Compounds 1. Predict the location of the atoms - • Hydrogen is always a terminal or end atom; it only shares 1 electron pair. • The atom with the least attraction for shared electrons is the central atom - the central atom is usually the one closer to the left in the periodic table. • All other atoms become terminal atoms. • For example, N is the central atom in NH3; C is the central atom in CO2.

5. Determining Lewis Structures of Molecular Compounds 2.Find the total number of electrons available for bonding by adding up the number of valence electrons of each atom in the compound. • In NH3, N has 5 and each H has 1, so the total is 8. • In CO2, C has 4 and each O has 6, so the total is 16.

6. Determining Lewis Structures of Molecular Compounds 3. Determine the number of bonding pairs by dividing the total number of electrons available by 2. • For NH3: 8/2 = 4 • For CO2: 16/2 = 8

7. Determining Lewis Structures of Molecular Compounds 4. Place one bonding pair between the central atom and each terminal atom. H N H C O O H

8. Determining Lewis Structures of Molecular Compounds 5A. Subtract the number of pairs you used in step 4 from the total pairs you determined in step 3. • For NH3, there are 4 bonding pairs and 3 were used in step 4. One bonding pair remains. • For CO2, there are 8 bonding pairs and 2 were used in step 4. 6 bonding pairs remain.

9. Determining Lewis Structures of Molecular Compounds 5B. The remaining pairs include lone pairs and pairs used in double and triple bonds. • First, place lone pairs around each terminal atom (EXCEPT HYDROGEN) to satisfy the octet rule. • Next, assign any remaining pairs to the central atom.

10. In NH3, 1 bonding pair remained. Lone pairs cannot be placed around H, so it must go on the N. H – N – H H In CO2, 6 bonding pairs remained, so they are placed around the O’s. O – C - O

11. Determining Lewis Structures of Molecular Compounds 6. Determine whether the central atom satisfies the octet rule. If it doesn’t, you MUST convert one or two of the lone pairs on the terminal atoms into a double or triple bond between the terminal and central atoms. Keep in mind, carbon, nitrogen, oxygen and sulfur often form double and triple bonds.

12. O – C - O Since the C in CO2 is only surrounded by 2 electron pairs, a lone pair on each O must be converted to a double bond. O = C = O

13. Practice Problems • NF3 • CS2 • H2Se

14. Polyatomic ions • The atoms within the ion are covalently bonded. • Drawing Lewis structures is similar to the procedure just practiced. The difference is in step #2. • Determine the number of electrons available in the atoms of the ion • Then subtract the ion charge if it is positive and add the ion charge if it is negative • Put a bracket around the ion structure with the charge as a superscript

15. Example: PO4-3 • The number of valence electrons available are 29 • 3 must be added to this • Total of 32 electrons are available - this is 16 bonding pairs

16. O 0 – P – 0 O 3-

17. Practice Problems • ClO4- • NH4+ • AsO4-3 • TeO4-2

18. Resonance Structures • Using the same sequence of atoms, it is possible to have more than one correct Lewis structure when a molecule or polyatomic ion has both a double bond and a single bond.

20. Resonance Structures • Resonance is a condition that occurs when more than one Lewis structure can be written for a molecule or ion. • Resonance structures are the 2 or more correct Lewis structures for a molecule or ion • They differ ONLY in position of the lone & bonding pairs, NEVER atom positions

21. Resonance Structures • Each actual molecule or ion that undergoes resonance behaves as if it has only one structure. • The actual bond lengths are equal to each other - the bond length is an average of the bonds in the resonance structure. • All bonds are shorter than single bonds but longer than double bonds!

22. Practice Problems • Draw the resonance structures of • O3 • NO2- • SO4-2

23. Exceptions to the Octet Rule • Some molecules and ions do not obey the octet rule. • Three (3) reasons exist for these exceptions. • First, a small group of molecules has an odd number of valence electrons and cannot form an octet around each atom.

24. Exceptions to the Octet Rule

25. Exceptions to the Octet Rule • Second, some compounds form with fewer than 8 electrons present around an atom – a suboctet. • This group is relatively rare; BH3 is an example.

26. Compounds like BH3 tend to be reactive - they will share an entire pair donated by another atom. • If this occurs, the bond formed is called a coordinate covalent bond. • Definition: a bond formed when an atom donates a pair of electrons to be shared with an atom/ion that needs 2 electrons to become stable. • Atoms/ions with lone pairs often form coordinate covalent bonds with atoms/ions that need 2 electrons

28. Exceptions to the Octet Rule • Third, some compounds have central atoms that contain more than 8 valence electrons. • This electron arrangement is referred to as an expanded octet. • Expanded octets are explained by considering the d orbitals of certain elements • Lewis structures are drawn by adding lone pairs or more than 4 bonding pair to the central atom

29. Exceptions to the Octet Rule