1. Basic Concepts of Chemical Bonding Chapter 8

2. Three Types of Chemical Bonds Ionic bond Ionic bond –Transfer of electrons –Between metal and nonmetal ions Metallic bond Metallic bond –Bonding electrons relatively free to move Covalent bond Covalent bond –Sharing of electrons –Between nonmetal atoms

3. H F F H Polar covalent bond ≡ a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ -- Electronegativity: the ability of an atom in a molecule to attract electrons to itself

4. Fig 8.6 Electronegativities of the Elements

5. Figure 8.7

6. Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond. Most polar Least polar Table 8.3 Fig 8.9

7. Add 1 for each negative charge. Subtract 1 for each positive charge. 1.Sum up all valence electrons. Add 1 for each negative charge. Subtract 1 for each positive charge. 2.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 3.Complete octets of atoms connected to central atom. 4.Place remaining electrons on central atom. 5.If not enough electrons to give central atom an octet, try multiple bonds. Writing Lewis Structures (p 314)

8. Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 2 – N is less electronegative than F, put N in center FNF F Step 1 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9. Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 2 – C is less electronegative than O, put C in center OCO O Step 1 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e - 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

10. Formal Charges Formal charge - the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons -

11. The best Lewis structure…  …is the one with the fewest charges  …puts a negative charge on the most electronegative atom. Formal Charges

12. Draw Lewis structure for ozone, O 3 or

13. But this is at odds with the true, observed structure of ozone, in which… …both O−O bonds are the same length:

14. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

15. Resonance structure - one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. OOOOOO OCO O OCO O OCO O What are the resonance structures of the carbonate (CO 3 2- ) ion? e.g., ozone 2-

16. Resonance The organic compound benzene, C 6 H 6, has two resonance structures: It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

17. Exceptions to the Octet Rule Too few electrons HHBe Be – 2e - 2H – (2)1e - 4e - BeH 2 BF 3 B – 3e - 3F – (3)7e - 24e - FBF F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24

18. Exceptions to the Octet Rule Odd number of electrons N – 5e - O – 6e - 11e - NO N O Too many electrons (central atom with principal quantum number n > 2) SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48

19. Average Bond Enthalpies Bond Type Bond Enthalpy (kJ/mol) C‒CC‒C 348 C=CC=C 614 C≡CC≡C 839 C‒NC‒N 293 C=NC=N 615 C≡NC≡N 891 Bond Enthalpies Single bond < Double Bond < Triple Bond

20. Chemical Bonding  These are average bond enthalpies, not absolute bond enthalpies  The C−H bonds in methane, CH 4, will be a bit different than the C−H bond in chloroform, CHCl 3 Table 8.4 Average bond Enthalpies (kJ/mol)

21. Fig 8.14 Estimating Enthalpies of Reaction  H rxn =  (bond enthalpies of bonds broken) -  (bond enthalpies of bonds formed)

22. In this example: one C-H bond and one Cl- Cl bond are broken one C-Cl and one H-Cl bond are formed Fig 8.14 Estimating Enthalpies of Reaction CH 4 (g) + Cl 2 (g)  CH 3 Cl (g) + HCl (g)

23. So,  H = [D(C−H) + D(Cl−Cl)] − [D(C−Cl) + D(H−Cl)] = [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)] = (655 kJ) - (759 kJ) = -104 kJ Fig 8.14 Estimating Enthalpies of Reaction CH 4 (g) + Cl 2 (g)  CH 3 Cl (g) + HCl (g)

24. Bond Type Bond Length (pm) C‒CC‒C 154 C=CC=C 133 C≡CC≡C 120 C‒NC‒N 143 C=NC=N 138 C≡NC≡N 116 Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond