Chapter 8 Covalent Bonding

1. Chapter 8Covalent Bonding Sec. 8.1: The Covalent Bond

2. The Covalent Bond • Objectives • Apply the octet rule to atoms that bond covalently. • Describe the formation of single, double, and triple covalent bonds. • Compare & contrast sigma & pi bonds • Relate the strength of covalent bonds to bond length and bond dissociation energy

3. Most compounds, including those in living organisms are covalently bonded.

4. Why do atoms bond? • The stability of a substance is related to its energy: lower energy states are more stable. • Metals and nonmetals gain stability and a lower energy state by transferring electrons. The ions that form have stable noble-gas configurations. • (Ions of opposite charge are then attracted to each other in an ionic bond.)

5. According to the octet rule, there is another way atoms can gain stability. They can SHARE valence electrons to achieve a noble-gas configuration.

6. The Covalent Bond • The chemical bonds that results from the sharing of valence electrons is a covalent bond.

7. The Covalent Bond • In a covalent bond, the shared electrons are considered to be part of the complete outer energy level of both atoms involved. • Covalent bonding occurs between NONMETAL ATOMS. • The elements are generally relatively close to each other on the periodic table.

8. What is happening in that covalent bond?? The nucleus of 1 atom is attracted to the electrons of another atom and vice versa.

9. The Covalent Bond • A distance between 2 atoms is reached where the attraction-repulsion forces are balanced. • This is a point of maximum stability for the atoms. • This is the point of covalent bond formation.

10. The Covalent Bond • A molecule is formed when two or more atoms bond covalently. • Compounds with covalent bonds are called covalent or molecularcompounds.

11. The Covalent Bond • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, and Iodine occur in nature as diatomic molecules because these 2-atom molecules are more stable than single atoms.

12. The Covalent Bond • When a single pair of electrons is shared, a single covalent bond is formed. H2 has a single covalent bond.

13. The Single Covalent Bond • The diatomic molecules of Group 7A elements also have single covalent bonds because they only need 1 more electron to attain noble gas configurations. • Example: • Draw the Lewis dot structure of F2 • Note that it has 3 unshared pairs (also called lone pairs) and 1 shared (or bonding) pair of valence electrons. • In a diagram, the shared pair can be represented by a single dash

14. What about oxygen? • Each O atom has 2 unshared pairs of valence electrons. • Therefore, O2 has 2 bonding pairs. • O2 has a double covalent bond. • Double bonds are represented in diagrams by a double dash: = O = O

15. Nitrogen? • 3 shared pairs • a triple covalent bond • Each atom has one unshared or lone pair ..

16. Water • Recall oxygen and all group 6A atoms have 6 valence electrons. • There are 2 electrons available to form covalent bonds. • Oxygen will share 1 with each hydrogen. • 2 single covalent bonds form - 2 lone pairs remain

17. Ammonia • Recall nitrogen and all group 5A atoms have 5 valence electrons. • N has 3 electrons that are available to form covalent bonds. • Nitrogen will share one with each hydrogen. • 3 single covalent bonds form with one lone pair

18. Group 4A • Will form 4 covalent bonds • For example, methane (CH4) forms 4 C-H covalent bonds

19. The Covalent Bond • Draw the Lewis structure for each of these molecules: • PH3 • H2S • CCl4

20. Sigma (σ) Bonds • Single covalent bonds are also called sigma bonds • The shared electron pair is in an area centered between the 2 atoms • Atomic orbitals overlap and merge • “s” with another “s” • “s” with a “p” • “p” with another “p”

21. A new, hybridized orbital is formed. It is called a bonding orbital – it is a localized region where bonding electrons will most likely be found. Illustrated are the sigma bonds in methane (text p.245).

22. Multiple Covalent bonds • A miltiple bond consists of one sigma (σ) bond and at least one pi (π) bond. • The pi bondis formed when parallel orbitals overlap and share electrons. • The pi bond occupies the space above and below the line that represents where the two atoms are joined together.

23. The Pi bond in Ethene (C2H4)

24. Strength of covalent bonds • Strength depends on how much distance there is between the nuclei of bonded atoms. • BOND LENGTH is the distance from the center of one nucleus to the center of the other nucleus of 2 bonded atoms at the point of maximum attraction.

25. Strength of covalent bonds • The shorter the bond length, the stronger the bond • As the number of shared pairs increases, the bond length decreases, so bond strength increases. • the bonds in F2 are weaker than those in O2 • the bonds in O2 are weaker than those in N2

26. Energy is released when bonds form • The bond dissociation energy is the energy added to break bonds. • It is “+” in value • Bond strength is described in terms of bond dissociation energy. • for F2, BDE equals 159 kJ/mol • for N2, BDE equals 945 kJ/mol; N2 has more shared pairs, so it has a shorter and stronger bond!

27. The sum of the BDE’s for all the bonds in a compound give an indication of the potential energy available in one molecule of the compound.

28. Chemical Reactions(an overall view) • In all chemical reactions, 2 changes in bonds must occur • Bonds between reactants MUST break; requires energy. • Bonds between products MUST form; releases energy. • Determining whether a reaction is endothermic or exothermic depends on the net energy change of the reaction.

29. A reaction is EXOTHERMIC if more energy is released when new bonds form than is required to break bonds. A reaction is ENDOTHERMIC if more energy is required to break bonds than is released when new bonds form. Chemical Reactions Either way, the net energy change is called the HEAT OF REACTION.