Covalent Bonding

1. Chemistry 11 Resource: Chang’s Chemistry Chapter 9 Covalent Bonding

2. Objectives • Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei. • Describe how the covalent bond is formed as a result of electron sharing. • Deduce the Lewis structures of molecules and ions for up to four electron pairs on each atom. • State and explain the relationship between the number of bonds, bond length and bond strength. • Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values.

3. Objectives • Predict the relative polarity of bonds from electronegativity values. • Predict the shape and bond angles for species with four, three, and two negative charge centers on the central atom using the valence shell electron pair repulsion theory (VSEPR). • Predict whether or not a molecule is polar from its molecular shape and bond polarities.

4. The covalent bond • A bond in which two electrons are shared by two atoms. • Covalent compounds are compounds that contain only covalent bonds. • Most of the molecules found in nature contain covalent bonds. • Model

5. The covalent bond Ionic vs. covalent bonding Why do atoms form covalent bonds?

6. The covalent bond • Covalent bonds are formed to achieve stability—the octet rule. • An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons (or the configuration of a noble gas). • Exceptions to the octet rule

7. The covalent bond Ways of representing the covalent bond Cl - Cl

8. The covalent bond Draw the Lewis structures for: • H2 • CH4

9. The covalent bond H – H Hi H – C – H Hi

10. The covalent bond Draw the Lewis structure for ammonia (NH3)

11. The covalent bond Ammonia (NH3)

12. The covalent bond Draw the Lewis structures for: • H2O • PCl3

13. The covalent bond lone pair

14. Multiple covalent bonds Number of electrons shared

15. Multiple covalent bonds Representations of a double covalent bond O = O

16. Multiple covalent bonds How do multiple covalent bonds affect the atoms that are bonded together? Is a single bond stronger or weaker than a double bond?

17. Multiple covalent bonds Bond energies of some diatomic molecules

18. Multiple covalent bonds Average bond lengths of common covalent bonds

19. Multiple covalent bonds The greater the number of covalent bonds, the stronger and shorter the bond between two atoms.

20. Electronegativity • Recall the definition of electronegativity. • In covalent bonds, we expect the electrons to be equally shared. • Consider the covalently bonded compound HF (p 357 Figure 9.4) • Why do the electrons “spend more time” around the fluorine nucleus?

21. Electronegativity • We can predict the relative electronegativity of an element using the periodic table based on the following factors: • shielding • nuclear charge

22. Electronegativity • This difference in electronegativity results in a polar covalent bond.

23. Electronegativity So what’s the difference between a polar covalent bond and an ionic bond?

24. Electronegativity • The 2.0 rule of electronegativity: A difference in electronegativity greater than or equal to 2.0 will most likely indicate an ionic bond. How else can we predict whether or not a bond between two elements is ionic or covalent?

25. Electronegativity Classify the following bonds as ionic, polar covalent, or covalent: • the bond in hydrochloric acid (HCl) • the bond in potassium fluoride (KF) • the CC bond in ethane (H3CCH3)

26. Electronegativity • the bond in HCl – polar covalent • the bond in KF - ionic • the CC bond in ethane (H3CCH3) – purely covalent Typically, only atoms of the same element can be joined by a pure covalent bond.

27. Practice exercises Covalent bonding: p 379, #9.30-9.32 Electronegativity and bond type p 379 #9.36, 38, 40 Lewis structures p 379 #9.44