Molecular Structure and Covalent Bonding Theories

1. Molecular Structure and Covalent Bonding Theories Chapter 8

2. The Valence Shell Electrons • Valence shell electrons • These electrons are largely responsible for____ • Electrons not present in the preceding ___ ___ • Ignore filled sets of d and f orbitals • Used to determine the Lewis structure of a compound containing covalent bonds • Works well for molecules containing atoms from the ____ ____ elements

3. Models to Describe Covalent Bonding • Valence shell electron pair repulsion (VSEPR) model – predicts the _____ ______ of atoms in a molecule • This will be related to a physical property called ______ • Valence bond theory – predicts how bonding will take place by ______ of atomic orbitals

4. VSEPR Theory • Valence shell electron are present as either ___ ___ or ____ ____. • Regions of high electron density are created. • These regions arrange themselves to be as far away as possible form on another. As a result specific geometries are created around atoms in the molecule • Single, double, and triple bonds are counted as one region of electron density • Unshared pairs of valence electrons are also counted as one region of electron density Drawing the Lewis structure accurately will reveal the number of electron density regions around the center atoms

5. VSEPR Theory • Draw the Lewis structures for CO2, H2CO, and CH4 • How will these regions of electron density arrange themselves to be as far away as possible from one another? • There are five basic shapes based on the number of electron density regions around a center atom(s) • Illustration of models with next few slides

6. VSEPR Theory Two regions of high electron density

7. VSEPR Theory Three regions of high electron density

8. VSEPR Theory Four regions of high electron density

9. VSEPR Theory Five regions of high electron density

10. VSEPR Theory Six regions of high electron density

11. VSEPR Theory • Encountered geometries • Electronic geometry – determined by the location of “___” the regions of electron density around the center atom(s) • Molecular geometry – determined by the arrangement of _____ only around the center atom(s) • The does not include lone electron pairs. The molecular shape differs from the electron shape if lone pairs are present An example is H2O

12. VSEPR Theory • Lone pairs of electrons occupy more space than bonding pairs. As a consequence, there is an order of the magnitude of repulsions • lp/lp > lp/bp > bp/bp As a result, the bond angles around a center atoms can be distorted (reduced) from the predicted values CH4 and H2O What are the H-C-H and H-O-H bond angles. If a change is observed, why?

13. Molecular Geometry and Polarity • The polarity can be determine once the geometry is known • A polar bond is created if the atoms sharing the electron pair have different electronegativities • HCl and the associated dipole moment. This molecule is polar. For diatomics, determination of polarity is easy. What if the molecule has two or more atoms? All the dipole have to be summed. If the sum equals zero, the molecule has no dipole.

14. Molecular Geometry and Polarity • A dipole moment (bond dipole) has _____ and _____. Both must be considered when determines if a molecule is polar. • CO2 and H2O. Do these molecules have net dipoles? • Conditions for polarity • There must be at least one polar bond or lone pair on a central atom • The bond dipoles must not cancel or if there are two or more lone pairs on the central atom, they must not be arranged so that their polarities cancel CO2, H2O, and O3

15. Molecular Geometry and Polarity

16. Valence Bond(VB) Theory • VB theory describes how bonding occurs • Describes how the atomic orbitals overlap to produce the bonding geometry predicted by VSEPR • Go back and review atomic orbitals if necessary • Electrons are arranged in atomic orbitals according to energy. The set of atomic orbitals, however, may not be of lowest possible energy upon bonding covalently to neighboring atoms.

17. Valence Bond(VB) Theory • The valence shell orbitals (atomic orbitals) commonly combine to change their character in order to obtain a lower energy ‘mixed’ orbital set for bonding in a particular geometry • Which atomic orbitals would participate in bonding in H, O, and C? These atomic orbitals can form a new set of hybrid orbitals upon bonding. • Hybrization – process by which ____ ____ combine to form a set of ‘mixed’ orbitals of lower energy when bonding covalently • The ‘mixed’ orbitals are called hybrid orbitals

18. Valence Bond(VB) Theory • Hybrid orbitals on a center atom align themselves with the bonding orbitals on the neighboring atoms • A ‘good overlap’ is necessary for sharing electrons in a bond. • Table 8-2 (refer to it) • The label given to a set of hybridized orbitals reflects the number and type of atomic orbitals used to produce the set. • Indicates the electronic geometry in agreement with VSEPR

19. Valence Bond(VB) Theory

20. Molecular Shapes and Bonding • Simples structures will be analyzed based on geometry type. • Experimentally determined findings will be discussed in light of these models. • Terminology • A – central atom • B – atoms bonded to A • U – lone pairs of electrons around A AB3U represents three atoms bonded to a central atom with one lone pair. An example would be NH3

21. Molecular Shapes and Bonding Discussion sequence • Experimental facts and Lewis formula • VSEPR • Electronic geometry • Molecular geometry • Polarity • Valence bond theory

22. AB2 Molecules - No Lone Pairs on A - Linear Molecules • The BeCl2 molecule is linear and has melting point of 405C. • Draw BeCl2 and discuss electronic geometry • Does the molecular geometry differ? • The molecule does not satisfy the octet rule • The compound bonds covalently due to the high charge density on Be2+ • The electron cloud on the halide is distorted by the high charge density • BeBr2 and BeI2 also have linear geometries

23. AB2 Molecules - No Lone Pairs on A - Linear Molecules • The molecule possesses two polar bonds (Be-Cl) • EN = 1.5 • The molecule, however, has no net dipole because the two bond dipoles are equal but in opposite directions. : : : : Cl-Be-Cl : : Bond dipoles cancel. This is a nonpolar molecule.

24. AB2 Molecules - No Lone Pairs on A - Linear Molecules • Electronic StructuresLewis Formulas 1s2s2p Be ­¯­¯ 3s3p Cl [Ne] ­¯­¯­¯­ The 2s orbital is full indicating that it will not bond. How will the Be atom make these electrons available for bonding? What happens in this molecule? Experimental data indicates that the Be-Cl bonds are identical.

25. AB2 Molecules - No Lone Pairs on A - Linear Molecules • Valence Bond Theory (Hybridization) 1s2s2p1ssp hyb2p Be ­¯ ­¯ Þ­¯ ­ ­ 3s3p Cl [Ne] ­¯­¯­¯­ The two atomic orbitals on Be hybridize to produce two sp hybrid orbitals that have properties between the s and p atomic orbitals. Notice that chlorine has a half-filled 3p orbital that can overlap with the sp hybrid orbitals of Be.

26. AB2 Molecules - No Lone Pairs on A - Linear Molecules Two regions of electron density around the central atom Illustrate how the sp orbitals overlap with the 3p orbitals on Cl

27. AB3 Molecules - No Lone Pairs on A - Trigonal Planar Molecules • Group IIIA elements that form covalent compounds by bonding to three other atoms • Octet rule is not satisfied but no big deal • Boron trichloride is a trigonal molecule with a melting point of -107C • Does the molecular and electronic geometry differ? • The data indicates that this molecule is nonpolar (no net dipole).

28. AB3 Molecules - No Lone Pairs on A - Trigonal Planar Molecules • Lewis structure predicts trigonal planar geometry • There are three bond dipoles of equal length but different direction. • The bond dipoles cancel each other The molecule has no net dipole How about BCl2H? Cl  BCl  Cl

29. AB3 Molecules - No Lone Pairs on A - Trigonal Planar Molecules • Electronic StructuresLewis Formulas 1s2s2p B ­¯­¯­ 3s3p Cl [Ne] ­¯­¯­¯­ Suppose that an electron in the 2s atomic orbital is promoted to an empty 2p atomic orbital allowing for 3 unfilled atomic orbitals for bonding. This would produce, however, unequal energies for the three B-Cl bonds. . B: : . Cl: :

30. AB3 Molecules - No Lone Pairs on A - Trigonal Planar Molecules • Valence Bond Theory (Hybridization) 1s2s2p1ssp2 hybrid B ­¯ ­¯­ Þ ­¯­ ­ ­ 3s3p Cl [Ne]­¯­¯­¯­ The 2s and 2p atomic orbitals on B hybridize to produce three sp orbitals (sp2 hybrid). Notice that chlorine has a half-filled 3p orbital that can overlap with the sp2 hybrid orbitals of Be.

31. AB3 Molecules - No Lone Pairs on A - Trigonal Planar Molecules Three regions of electron density around the central atom Illustrate bonding with the Cl atoms on the hybridized B

32. AB4 Molecules - No Lone Pairs on A - Tetrahedral Molecules • Group IVA elements that form covalent compounds by bonding to four other atoms • Four electrons are shared and the octet rule is generally satisfied • CH4, methane, possesses a tetrahedral geometry and has a melting point of -182C • Would the molecular and electronic geometry differ? • The data indicate that the molecule is nonpolar.

33. AB4 Molecules - No Lone Pairs on A - Tetrahedral Molecules • Lewis structure predicts tetrahedral geometry • There are four small bond dipoles which cancel • The molecule is nonpolar • What about CCl3H and CH3Cl? • When the symmetry lowers, the molecule becomes polar. • Other molecules? CH4

34. AB4 Molecules - No Lone Pairs on A - Tetrahedral Molecules • Electronic StructuresLewis Formulas 2s2p C [He] ­¯­ ­ 1s H ­ Suppose that an electron in the 2s atomic orbital is promoted to an empty 2p atomic orbital allowing for 4 unfilled atomic orbitals for bonding. This would produce, however, unequal energies for the four C-H bonds. . C: . H .

35. AB4 Molecules - No Lone Pairs on A - Tetrahedral Molecules • Valence Bond 2s2pfour sp3 hybrid orbitals C [He] ­¯­ ­ Þ C [He]­ ­ ­ ­ 1s H ­ The 2s and 2p atomic orbitals on C hybridize to produce four sp orbitals (sp3 hybrid). Notice that hyrogen has a half-filled 1s orbital that can overlap with the sp3 hybrid orbitals of C. Many AB4 type molecules have this hybridization NH4+ is an AB4 type polyatomic ion

36. AB4 Molecules - No Lone Pairs on A - Tetrahedral Molecules Four regions of electron density around the central atom Illustrate how the hydrogen atomic orbitals bond to the hybridized carbon sp3 orbitals

37. Alkanes CnH2n+2 • alkanes are saturated hydrocarbons • have the general formula CnH2n+2. CH4 - methane C2H6 or (H3C-CH3) - ethane C3H8 or(H3C-CH2-CH3) - propane • C atoms are located at the center of a tetrahedron each alkane is a chain of interlocking tetrahedra C atom at the center of each tetrahedron enough H to form a total of four bonds for each C

38. AB3U Molecules - One Lone Pair - Pyramidal Molecules • Group VA elements (e.g. N) have five electrons in the valence and commonly bond to three atoms leaving a lone pair. • The octet rule is satisfied • The most common molecule is NH3. • How many regions of electron density around nitrogen? The bong angle is in this molecule is ~107. Why? • Other common molecules are NF3, PF3, and the polyatomic ion SO32-.

39. AB3U Molecules - One Lone Pair - Pyramidal Molecules • The Lewis structure predicts tetrahedral electronic geometry. • Is the molecular geometry different? • There are three bond dipoles? Detail. • Is the molecule polar? • How about NF3? How do the polarities of the two molecules compare (later)? :

40. AB3U Molecules - One Lone Pair - Pyramidal Molecules Electronic StructuresLewis Formulas 2s2p N [He] ­¯­ ­ ­ 2s2p F [He] ­¯­¯­¯­ 1s H ­ There are three half-filled atomic orbitals on the nitrogen (2p). The data suggests, however, that there are four nearly equivalent orbitals (not three). Three orbitals are for bonding and one for a lone pair.

41. AB3U Molecules - One Lone Pair - Pyramidal Molecules • Valence Bond 2s2p four sp3 hybrids N [He] ­¯­ ­ ­ Þ ­¯­ ­ ­ The 2s and 2p atomic orbitals hybridize to form four sp3 hybrid orbitals. This hybridization is also necessary to produce the correct geometry for bonding. Illustrate bonding with hydrogen 1s atomic orbital. Once again there are four regions of electron density around the center atom

42. AB3U Molecules - One Lone Pair - Pyramidal Molecules • Let’s compare NH3 with NF3. • The geometry of the both molecules is already known • Electronic geometry is _________ • Molecular geometry is _________ • How does the lone pair influence polarity? It’s contribution has to be included to determine polarity of a molecule.

43. AB3U Molecules - One Lone Pair - Pyramidal Molecules : • The bond dipoles go opposite directions on NH3 and NF3 • For NH3, the net dipole is enhanced by the lone pair. • For NF3, the net dipole is decreased due to the lone pair • Additionally, the H-N-H angle is greater than the F-N-F angle due to closer approach of the lone pair to nitrogen on NF3 H-N-H = 107.3 : F-N-F = 102.1

44. AB2U2 - Two Lone Pairs - V-Shaped Molecules • Group VIA elements (e.g. O) have six electrons in the valence and commonly bond to two atoms leaving two lone pairs. • The octet rule is generally satisfied • H2O is the most common molecule of this type. • The molecular geometry is ______ and the electronic geometry is _____ • Other examples of this type of molecule is H2S and OCl2

45. AB2U2 - Two Lone Pairs - V-Shaped Molecules • The Lewis structure predicts that the molecule is bent in agreement with experimental data. • The actual H-O-H bond angle is 104.5 due to repulsions from two lone pairs • There are two bond dipoles (O-H). Additionally, the net dipole is enhanced by the lone pairs. • Illustrate : :

46. : . O: . AB2U2 - Two Lone Pairs - V-Shaped Molecules • Electronic StructuresLewis Formulas 2s2p O [He] ­¯­¯­ ­ 1s H ­ There are two half-filled atomic orbitals on the nitrogen (2p). The data suggests, however, that there are four nearly equivalent orbitals (not two). Two orbitals are for bonding and two for lone pairs. H .

47. AB2U2 - Two Lone Pairs - V-Shaped Molecules • Valence Bond 2s2p four sp3 hybrids O [He] ­¯­¯ ­ ­ Þ ­¯­¯ ­ ­ The hybrid orbitals that are full belong to the lone pairs. The half-filled orbitals are used for bonding.

48. Trigonal Bipyramidal Electronic Geometry • AB5, AB4U, AB3U2, and AB2U3 • Hybridization is sp3d. • The lone pairs (if present) will arrange themselves to minimize repulsive forces. • lp/lp >> lp/bp > bp/bp • This geometry is common for P, As, and Sb. • All five valence electrons are shared (PF5)

49. Trigonal Bipyramidal Electronic Geometry, AB5 • The VSEPR theory predicts trigonal bipyramidal for the electronic and molecular geometry. • There are three equatorial atoms and to axial atoms. What are the bond angles? Are the individual bonds polar? Is the molecule (type AB5) polar? • Show molecule with this geometry.

50. Trigonal Bipyramidal Electronic Geometry, AB5 • Electronic StructuresLewis Formulas 3s3p P [Ne] ­¯­ ­ ­ 2s2p F [He] ­¯­¯­¯­ The 3d subshell is empty and participates in the rehybridization (sp3d).