Shape of Molecules & Chemical Bonding

1. Chemistry 1A General Chemistry Ch. 10: The Shape of Molecules Instructor: Dr. Orlando E. RaolaSanta Rosa Junior College

2. Chemical bonding

3. 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?

4. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

5. Fundamental Properties of Models • A model does not equal reality. • Models are oversimplifications, and are therefore often wrong. • Models become more complicated as they age. • We must understand the underlying assumptions in a model so that we don’t misuse it.

6. Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

7. Structure and bonding NN triple bond. Molecule is unreactive Phosphorus is a tetrahedron of P atoms. Very reactive! Red phosphorus, a polymer. Used in matches.

8. Forms of chemical bonds • There are 2 extreme forms of connecting or bonding atoms: • Ionic—complete transfer of 1 or more electrons from one atom to another • Covalent—some valence electrons shared between atoms • Most bonds are somewhere in between.

9. Forms of chemical bonds

10. Ionic Compounds Metal of low IE Nonmetal of high EA 2 Na(s) + Cl2(g)  2 Na+ + 2 Cl-

11. Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons.Electron sharingresults. Bond is a balance of attractive and repulsive forces.

12. •• •• Cl H H Cl • • + • • •• •• Bond formation A bond can result from a “head-to-head”overlapof atomic orbitals on neighboring atoms. Overlap of H (1s) and Cl (2p) Note that each atom has a single, unpaired electron.

13. Chemical Bonding: Objectives Objectivesare to understand: 1. valence e- distribution in molecules and ions. 2. molecular structures 3. bond properties and their effect on molecular properties.

14. G. N. Lewis 1875 - 1946 Electron Distribution in Molecules • Electron distribution is depicted withLewis electron dot structures • Valence electrons are distributed as shared orBOND PAIRS and unshared orLONE PAIRS.

15. Place one dot per valence electron on each of the four sides of the element symbol. Pair the dots (electrons) until all of the valence electrons are used. . . . : . . . : N . N . . N N : . . : . Lewis Electron-Dot Symbols For main group elements - Example: Nitrogen, N, has 5 valence electrons.

16. Lewis electron-dot symbols for elements in Periods 2 and 3 1 2 13 14 15 16 17 18

17. Figure 9.6 The Born-Haber cycle for lithium fluoride.

18. Energy balance in the formation of an ionic compound

19. Formation of an ionic solid • 1. Sublimation of the solid metal • M(s)  M(g) [endothermic] • 2. Ionization of the metal atoms • M(g)  M+(g) + e [endothermic] • 3. Dissociation of the nonmetal • 1/2X2(g)  X(g) [endothermic] • 4. Formation of X ions in the gas phase: • X(g) + e X(g) [exothermic] • 5. Formation of the solid MX • M+(g) + X(g)  MX(s) [quite exothermic]

20. Lattice energy calculations Q1, Q2 = charges on the ions r = shortest distance between centers of the cations and anions

21. Born-Haber cycle for MgO

22. Born-Haber cycle for MgO

23. Compare lattice energy of LiF and MgO Similar ionic radii: Mg2+:72 pm O2-: 140 pm Li+: 76 pm F-: 133 pm Hlattice MgO = 3929 kJ·mol-1 Hlattice LiF = 1050 kJ·mol-1

25. Trends in lattice energy.

26. Valence electrons are distributed as shared orBOND PAIRS and unshared orLONE PAIRS. •• H Cl • • •• lone pair (LP) shared or bond pair Bond and Lone Pairs This is called a LEWIS ELECTRON DOT structure.

27. Valence Electrons Electrons are divided between core and valence electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5

28. Number of valence electrons: Main group elements: (1,2,13-18): - total of s and p electrons in the outer shell Transition elements: (3-12): - ns and (n-1) d Valence electrons

29. Octet rule Each atom has a tendency to be surrounded by 8 electrons in molecules and polyatomic ions (H is surrounded by 2 electrons).

30. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. (H  X)expected = ½(H-H + X-X)  = (H  X)actual (H  X)expected

31. The Pauling electronegativity (EN) scale

32. Electronegativity and atomic size

33. Lewis Dot Diagrams • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration.

34. Rules for forming Lewis dot diagrams • Determine the total number of valence electrons in the molecule or ion • Determine the arrangement of atoms within a molecule (the least electronegative is central, halogens are central only with oxygen, oxygen is central only in water, hydrogen is always terminal. • Use a pair of electrons to form a bond between each pair of atoms • Arrange the remaining electrons as lone pairs to satisfy the duet rule for hydrogen or the octet rule for other terminal atoms • If the central atom has fewer than eight electrons, move lone pairs from the terminal atoms to form multiple bonds.

35. Comments About the Octet Rule • 2nd row elements C, N, O, F observe the octet rule. • 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. • 3rd row and heavier elements CANexceed the octet rule using empty valence d orbitals. • When writing Lewis structures, satisfy octetsfirst,then place electrons around elements having available d orbitals.

36. Writing Lewis structures nitrogren trifluoride NF3 • Number of valence electrons: 5+(3x7)=26 • Use bond pairs to form bonds between atoms • Arrange remaining electrons to satisfy octet rule (duet for H)

37. More Lewis structures hydrazine N2H4 chloroethane C2H5Cl phosgene COCl2 chlorate ion ClO3- perchlorate ion ClO4-

38. Ammonia, NH3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. Therefore, N is central 2. Count valence electrons H = 1 and N = 5 Total = (3  1) + 5 = 8 electrons / 4 pairs Building a dot structure

39. H H N H •• H H N H Building a dot structure 3. Form a single bond between the central atom and each surrounding atom 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 H shares 1 pair.

40. 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 = 26 e- or 13 pairs Step 3. Form bonds

41. •• O • • • • •• •• O S O • • • • •• •• Sulfite ion, SO32- Remaining pairs become lone pairs, first on outside atoms and then on central atom. •• Each atom is surrounded by an octet of electrons.

42. Carbon Dioxide, CO2 1. Central atom = _______ 2. Valence electrons = __ or __ pairs 3. Form bonds. This leaves 6 pairs. 4. Place lone pairs on outer atoms.

43. Carbon Dioxide, CO2 4. Place lone pairs on outer atoms. 5. So that C has an octet, we shall form DOUBLE BONDS between C and O. The second bonding pair forms api (π)bond.

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

45. OR bring in bring in right pair left pair •• •• •• O S O • • • • •• •• Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs 3. Form double bond so that S has an octet — but note that there are two ways of doing this.

46. Sulfur Dioxide, SO2 This leads to the following structures. These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is aHYBRIDof the two.

47. Urea, (NH2)2CO

48. Urea, (NH2)2CO 1. Number of valence electrons = 24 e- 2. Draw sigma bonds.

49. Urea, (NH2)2CO 3. Place remaining electron pairs in the molecule.

50. Urea, (NH2)2CO 4. Complete C atom octet with double bond.