Chemical Bonding and Molecular Structure (Chapter 9)

1. Chemical Bonding and Molecular Structure (Chapter 9) • Ionic vs. covalent bonding • Molecular orbitals and the covalent bond (Ch. 10) • Valence electron Lewis dot structures • octet vs. non-octet • resonance structures • formal charges • VSEPR - predicting shapes of molecules • Bond properties • polarity, bond order, bond strength Bonding and Structure

2. Chemical Bonding Problems and questions — • How is a molecule or polyatomic ion held together? • Why are atoms distributed at strange angles? • Why are molecules not flat? • Can we predict the structure? • How is structure related to chemical and physical properties? Bonding and Structure

3. Forms of Chemical Bonds • There are 2 extreme forms of connecting or bonding atoms: • Ionic—complete transfer of electrons from one atom to another • Covalent—electrons shared between atoms Most bonds are somewhere in between. Bonding and Structure

4. Ionic Bonds Ionic compounds - essentially complete electron transfer from an element of low IE (metal) to an element of high electron affinity (EA) (nonmetal) Na(s) + 1/2 Cl2(g)  Na+ + Cl-  NaCl (s) - primarily between metals (Grps 1A, 2A and transition metals) and nonmetals(esp O and halogens) - NON-DIRECTIONAL bonding via Coulomb (charge) interaction Bonding and Structure

5. Br Br Covalent Bonding Covalent bond is the sharing of the VALENCE ELECTRONS of each atom in a bond Recall: Electrons are divided between core and valence electrons. ATOM core valence Na 1s2 2s2 2p6 3s1 [Ne] 3s1 Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5 Bonding and Structure

6. 8A 1A Valence Electrons 2A 3A 4A 5A 6A 7A Number of valence electrons is equal to the Group number. Bonding and Structure

7. Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons. A covalent bond is a balance of attractive and repulsive forces. 6_H2bond.mov Bonding and Structure

8. •• •• Cl H H Cl • • + • • •• •• Overlap of H (1s) and Cl (2p) Bond Formation A bond can result from a “head-to-head” overlapof atomic orbitals on neighboring atoms. This type of overlap places bonding electrons in a MOLECULAR ORBITAL along the line between the two atoms and forms a SIGMA BOND (s). Bonding and Structure

9. 6_H2pot.mov Sigma Bond Formation by Orbital Overlap Two s Atomic Orbitals (A.O.s) overlap to form an s (sigma) Molecular Orbital (M.O.) Bonding and Structure

10. Sigma Bond Formation by Orbital Overlap Two s A.O.s overlap to from an s  M.O. Similarly, two p A.O.s can overlap end-on to from a p M.O. e.g. F2 Bonding and Structure

11. Electron Distribution in Molecules • Electron distribution is depicted with Lewis electron dot structures • Electrons are distributed as: • shared or BOND PAIRS and • unshared or LONE PAIRS. G. N. Lewis 1875 - 1946 Bonding and Structure

12. •• H Cl • • •• Unshared or lone pair (LP) shared or bond pair Bond and Lone Pairs • Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. This is a LEWIS ELECTRON DOT structure. Bonding and Structure

13. This observation is called the OCTET RULE Rules of Lewis Structures • For Groups 1A-4A (Li - C), • no. of BOND PAIRS = group number • No. of valence electrons of an atom = Group number • For Groups 5A-7A (N - F), • no. of BOND PAIRS = 8 - group No. • Except for H • (and atoms of 3rd and higher periods), • #Bond Pairs + #Lone Pairs = 4 Bonding and Structure

14. Building a Dot Structure 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. In ammonia, N is central Ammonia, NH3 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons or 4 pairs Bonding and Structure

15. •• H H N H Building a Dot Structure 3. Form a sigma bond between the central atom and surrounding atoms. 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while each H shares 1 pair. Bonding and Structure

16. Sulfite ion, SO32- Step 1. Central atom = S Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 6 + 18 + 2 = 26 e- or 13 pairs Step 3. Form sigma bonds 10 pairs of electrons are left. Bonding and Structure

17. •• O • • • • •• •• O S O • • • • •• •• •• Sulfite ion, SO32- (2) Remaining pairs become lone pairs, first on outside atoms then on central atom. Each atom is surrounded by an octet of electrons. NOTE - must add formal charges (O-, S+) for complete dot diagram Bonding and Structure

18. Carbon Dioxide, CO2 1. Central atom = __C____ 2. Valence electrons = _16_ or _8_ pairs 3. Form sigma bonds. This leaves __6__ pairs. 4. Place lone pairs on outer atoms. Bonding and Structure

19. Carbon Dioxide, CO2 (2) 4. Place lone pairs on outer atoms. 5. To give C an octet, form DOUBLE BONDS between C and O. The second bonding pair forms a pi (p)bond. Bonding and Structure

20. H2CO SO3 C2F4 Double and even triple bonds are commonly observed for C, N, P, O, and S Bonding and Structure

21. Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 6 + 2*6 = 18 electrons or 9 pairs 3. Form pi () bond so that S has an octet — note that there are two ways of doing this. Bonding and Structure

22. Equivalent structures called: Sulfur Dioxide, SO2 RESONANCE STRUCTURES The proper Lewis structure is a HYBRID of the two. A BETTER representation of SO2 is made by forming 2 double bonds Each atom has - OCTET - formal charge = 0 O = S = O Bonding and Structure

23. Urea (NH2)2CO 1. Number of valence electrons = 24 e- 2. Draw sigma bonds. Leaves 24 - 14 = 10 e- pairs. 3. Complete C atom octet with double bond. 4. Place remaining electron pairs on oxygen and nitrogen atoms. Bonding and Structure

24. BF3 SF4 Violations of the Octet Rule elements of higher periods. Usually occurs with: Boron Bonding and Structure

25. Boron Trifluoride • Central atom = B • Valence electrons = 3 + 3*7 = 24 or electron pairs = 12 • Assemble dot structure The B atom has a share in only 6 electrons (or 3 pairs). B atom in many molecules is electron deficient. Bonding and Structure

26. Sulfur Tetrafluoride, SF4 • Central atom = S • Valence electrons = 6 + 4*7 = 34 e- or 17 pairs. • Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period. Bonding and Structure

27. Formal Atom Charges Formal charge = Group no. - 1/2 (no. bond electrons) - (no. of LP electrons) • Atoms in molecules often bear a charge (+ or -). • The most important dominant resonance structure • of a molecule is the one with formal charges • as close to 0 as possible. Bonding and Structure

28. 6 - ( 1 / 2 ) ( 4 ) - 4 = 0 • • • • O O C • • • • 4 - ( 1 / 2 ) ( 8 ) - 0 = 0 Carbon Dioxide, CO2 At OXYGEN At CARBON Bonding and Structure

29. 6 - ( 1 / 2 ) ( 2 ) - 6 = - 1 • • O O C • • • • • • 6 - ( 1 / 2 ) ( 6 ) - 2 = + 1 + • • O O C • • • • • • Carbon Dioxide, CO2 (2) An alternate Lewis structure is: + C atom charge is 0. AND the corresponding resonance form Bonding and Structure

30. + + • • O O C • • • • • • • • • • O O C • • • • • • -0.73 -0.73 O O C • • • • +1.46 • • Carbon Dioxide, CO2 (3) Which is the predominant resonance structure? • REALITY: Partial charges calculated by CAChe molecular modeling system (on CD-ROM). OR Answer ? Form without formal charges is BETTER - no +ve charge on O Bonding and Structure

31. Boron Trifluoride, BF3 What if we form a B—F double bond to satisfy the B atom octet? Bonding and Structure

32. •• F • • •• F B • • •• F • • • • •• Boron Trifluoride, BF3 (2) + fc = 7 - 2 - 4 = +1 Fluorine • To have +1 charge on F, with its very high electron affinity is not good. -ve charges best placed on atoms with high EA. • Similarly -1 charge on B is bad • NOT important Lewis structure fc = 3 - 4 - 0 = -1 Boron Bonding and Structure

33. A. S=C=N Calculated partial charges B. S=C - N C. S-C N -0.16 -0.32 -0.52 Thiocyanate ion, (SCN)- Which of three possible resonance structures is most important? ANSWER: C > A > B Bonding and Structure