2. Molecular Geometry and Bonding Theories Brown, LeMay Ch 9 AP Chemistry Monta Vista High School Ch 9 PS #8, 10, 16, 20 (and draw overall dipole moments), 30, 34, 36, 41, 59, 73; Recommended #9, 11, 13, 19, 25, 31, 40, 55, 60

3. Lewis structure is a flat drawing showing the relative placement of atoms, bonds etc. in a molecule, but does not tell anything about the shape of the molecule. VSEPR theory helps construct molecular shape (3-D) from the Lewis structures, which are 2-D structures. Rationale for VSEPR Theory

4. Valence Shell Electron Pair Repulsion Theroy The basis principal of VSEPR is that each group of valence electrons (electron domains) around a central atom tend to be as far as possible from each other to minimize repulsions. These electron domain repulsions around the central atom determine the molecular geometry of a molecule. Electron domains: areas of valence e - density around the central atom;  Includes bonding e - pairs and nonbonding (lone) e - pairs  A single, double, or triple bond counts as one domain Repulsions between two e domains are : lone pair- lonepair >lone pair-bond pair> bond pair-bond pair

5. Valence-shell electron-pair repulsion theory – Why is lone pair- lone pair repulsion greater than bond pair-bond pair repulsion? – Good Links: Good Power point on VSEPR, Shapes of sp3, sp2, sp Carbon, VSEPR Lecture Good Power point on VSEPRShapes of sp3, sp2, sp CarbonVSEPR Lecture Summary of ABE (Tables 9.1 - 9.3) on the next slide: E = lone or non-bonding pairs A = central atom B = bonded atoms Bond angles notation used here: < xº means ~2-3º less than predicted << xº means ~4-6º less than predicted

6. Tables Link Link # of e - domains & # and type of hybrid orbitals e - domain geometry Formula & Molecular geometry Predicted bond angle(s) Example (Lewis structure with molecular shape) 2 Two sp hybrid orbitals Linear AB 2 Linear 180º BeF 2 CO 2 A |X|X X A |B|B B

7. 3 Three sp 2 hybrid orbitals Trigonal planar AB 3 Trigonal planar 120º BF 3 Cl-C-Cl << 120º Cl 2 CO LAB 2 Bent < 120º NO 2 1- A |X|X XX A |B|B BB A |B|B :B

8. Example: CH 4 Molecular shape = tetrahedral Bond angle = 109.5º H | H—C—H | H 109.5º

9. 4 Four sp 3 hybrid orbitals or Tetrahedral AB 4 Tetrahedral 109.5º CH 4 LAB 3 Trigonal pyramidal < 109.5º Ex: NH 3 = 107º NH 3 L 2 AB 2 Bent <<109.5º Ex: H 2 O = 104.5º H2OH2O X A X X X X A X XX B A B BB : A B BB : A B B:

10. PCl 5 Molecular shape = trigonal bipyramidal Bond angles equatorial = 120º axial = 90º :Cl: \ / :Cl—P—Cl: | :Cl: : : :: : : : 120º 90º

11. 5 Five sp 3 d hybrid orbitals Trigonal bipyramidal AB 5 Trigonal bipyramidal Equatorial = 120º Axial = 90º PCl 5 LAB 4 Seesaw Equatorial < 120º Axial < 90º SF 4 X X X A X |X|X B B B A B |B|B : B - A - B B B

12. 5 Five sp 3 d hybrid orbitals Trigonal bipyramidal L 2 AB 3 T-shaped Axial << 90º ClF 3 L 3 AB 2 Linear Axial = 180º XeF 2 X X X A X |X|X B : : A B |B|B : : : A B |B|B

13. 6 Six sp 3 d 2 hybrid orbitals or Octahedral AB 6 Octahedral 90º SF 6 LAB 5 Square pyramidal < 90º BrF 5 X X A X |X|X X X B B A B |B|B B B B B A B |.. B B A X |X|X XX X X

14. 6 Six sp 3 d 2 hybrid orbitals or L 2 AB 4 Square planar 90º XeF 4 or L 3 AB 3 T-shaped <90º KrCl 3 1- A B |B|B BB.. B B A |.. B B B B A.. |.... B A |B|B BB

15. Limitations of VSEPR Theory Even though the VSEPR model is useful in predicting the shapes of molecules, it does not differentiate between single, double and triple bonds and does not account for bond strengths.

16. 9.3: Molecular Polarity Two factors must be considered in the polarity of a molecule: 1.Are the individual bonds (joining different atoms in a molecule) polar? Ex. HCl vs. H 2. HCl is polar while H 2 is not. H-Cl: : : : :  +  - Bond polarity is most often represented by an arrow that points toward the  - (most EN atom), showing the shift in e - density. The dipole moment (  ) is a vector (i.e., has a specific direction) measuring the polarity of a bond which contains partial charges (Q) that are separated by a distance (r).  = Q r

17. 2. If individual bonds are polar, then do individual dipole moments cancel out or not. A molecule is polar if its centers of (+) and (-) charge do not cancel out- generally happens in a distorted molecule (molecules with lone pairs of e on the central atom.) How to determine if a molecule is polar? The sum of the bond dipole moments in a molecule determines the overall polarity of the molecule. 1.Draw the true molecular geometry (3D geometry). 2.Draw each bond dipole as an arrow 3.Add the vectors, and draw the overall dipole moment. If none, then  = 0. 4.Generally, a distorted molecule (with lone pairs on the central atom) will have a dipole. Exception: AB2E3 type

18. What is the big deal about polarity? The polarity of a molecule will tell you a lot about its solubility, boiling point, etc. when you compare it to other similar molecules. Water, for example, is a very light molecule (lighter than oxygen gas or nitrogen gas) and you might expect it would be a gas based on its molecular weight, however the polarity of water makes the molecules "stick together" very well. Hence water is present as liquid. Polar substances are soluble in water (which is polar) and non polar substances are soluble in non polar solvents such as benzene and oil.

19. Ex: Draw molecular geometires, bond dipole moments, and overall dipole moments. Also, name the e - domain geometry and the molecular geometry. CO 2 BF 3 H 2 O CCl 4 NH 3 PH 3

20. Rationale For Valence Bond Theory VB theory provides a basis for covalent bond formation based upon overlapping of atomic orbitals to share electrons. This theory successfully predicts bond strengths based upon orbital overlap (such as H 2 bond being weaker than N 2 bond)- sigma vs. pi bond strength

21. Covalent Bonding and Orbital Overlap: Valence Bond Theory The basic principle of VB theory is that a covalent bond forms when the orbitals of two atoms overlap. Three central themes of VB theory derive from this principle: 1. Opposing spins of e pairs: In accordance with Pauli’s exclusion principle, an orbital can have max of two e with opposite spins. 2. Maximum overlap of bonding orbitals: The bond strength depends upon the attraction of nuclei for the shared e, so the greater the overlap, the stronger the bond. 3. End to end overlap of the atomic orbitals form a sigma bond and allows the free rotation of the parts of the molecule. Side-to-side overlap forms a pi bond, which restricts rotation. A multiple bond consists of one sigma bond and rest pi bonds.

22. Sigma and Pi bonds Sigma (  ) bond: Covalent bond that results from axial overlap of orbitals between atoms in a molecule Lie directly on internuclear axis “Single” bonds, could form between s-s orbital or s-p orbital or p-p orbital by axial overlapping Ex: F 2 Pi (  ) bond: Covalent bond that results from side-by-side overlap of orbitals between atoms in a molecule. Are “above & below” and “left & right” of the inter nuclear axis and therefore have less total orbital overlap, so they are weaker than  bonds. Forms between two p orbitals (py or pz) Make up the 2 nd and 3 rd bonds in double & triple bonds. Ex: O 2 N 2

23. Covalent Bonding and Orbital Overlap Valence-bond theory: overlap of orbitals between atoms results in a shared valence e - pair (i.e., bonding pair) Energy (kJ/mol) 0 -436 0.74 Å H-H distance Figure : Formation of bond in H 2 a.As 2 H atoms approach, the 2 valence e- in the 1s orbitals begin to overlap, becoming more stable. b.As H-H distance approaches 0.74 Å, energy lowers b/c of electrostatic attraction between the nuclei & the incoming e-. c.When H-H distance = 0.74 Å, energy is at its lowest because electrostatic attractions & repulsions are balanced. (This is the actual H-H bond distance. d.When H-H distance < 0.74 Å, energy increases b/c of electrostatic repulsion between 2 nuclei & between the 2 e-. a b c d

24. Limitations of VB Theory Valence bond (VB) theory assumes that all bonds formed between two atoms are localized bonds and are formed by the donation of an electron from each atom. This is actually an invalid assumption because many atoms bond using delocalized electrons.

25. Rationale for Hybrid Orbital Theory Good youtube video Good youtube video Hybrid orbital theory is seen as an extension of VB theory, where atomic orbitals “hybridize” to form new hybrid orbitals. This hybrid orbital theory helps explains the bonding in terms of quantum mechanical model of atom (s,p,d,f orbitals).

26. 9.5: Hybrid Orbital Theory Movie on Hybrid Orbitals Movie on Hybrid Orbitals  Explains the molecular geometries in terms of s,p,d,f orbitals.  VSEPR explains that e domains must be farthest from each other around central atom, but fails to explain these in terms of orbitals as defined in wave mechanical model of atom.  Hybrid orbital theory of Linus Pauling proposed that the valence atomic orbitals in the molecule are very different from those in the isolated atoms.  The process of orbital mixing is called hybridization, and the new atomic orbitals are called hybrid orbitals. Animation on Hybrid Orbitals, Hybridization Movie Animation on Hybrid Orbitals Hybridization Movie

27. Hybrid Orbital Theory Two key points about the number and types of hybrid orbitals are that 1. The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. 2. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.

28. “sp” hybrid orbitals BeF 2 (g): observed as a linear molecule with 2 equal-length Be-F bonds. Valence bond theory predicts that each bond is an overlap of one Be 2s e - and one 2p e - of F. However, Be’s 2s e - are already paired. So… To form 2 equal bonds with 2 F atoms: 1.In Be, one 2s e - is promoted to an empty 2p orbital. 2.The occupied s and p orbitals are hybridized (“mixed”), producing two equivalent “sp” orbitals. 3.As the two “sp” hybrid orbitals of Be overlap with two p orbitals of F, stronger bonds result than would be expected from a normal Be s and F p overlap. (This makes up for energy needed to promote the Be e - originally.)

29. Be (ground state) → Be (promoted) → Be (sp hybrid) 2p 2s Energy → Orbital “shapes” One s + one p → Two sp orbitals (to bond with 2 F’s) A central atom in a Lewis structure with exactly 2 e - domains has sp hybrid orbitals. F

30. “sp 2 ” hybrid orbitals BF 3 (g): observed as trigonal planar molecule with 3 equal-length B-F bonds. However, 2 valence e- in B are paired, and are the s and p e - not at the observed 120 º angle. 2p 2s One s + two p → Three sp 2 orbitals (to bond with 3 F’s) A central atom with exactly 3 e- domains has sp 2 hybrid orbitals. B (ground) → B (promoted) → B (sp 2 hybrid) F F F

31. “sp 3 ” hybrid orbitals CH 4 (g): observed as tetrahedral 2p 2s One s + three p → Four sp 3 orbitals (to bond with 4 H’s) A central atom with exactly 4 e- domains has sp 3 hybrid orbitals. C (ground) → C (promoted) → C (sp 3 hybrid) H H

32. “sp 3 d” hybrid orbitals (or dsp 3 ) PCl 5 (g): observed as trigonal bipyramidal; forms 5 bonds of equal energy (* but not equal length: equatorial are slightly longer) 3d 3p 3s One s + three p + one d → Five sp 3 d orbitals (to bond with 5 Cl’s) A central atom with exactly 5 e- domains has sp 3 d hybrid orbitals. P (ground) → P (promoted) → P (sp 3 d hybrid) Cl Cl Cl Cl Cl

33. “sp 3 d 2 ” hybrid orbitals (or d 2 sp 3 ) SF 6 (g): observed as octahedral; forms 6 equal-length bonds One s + three p + two d → Six sp 3 d 2 orbitals A central atom with exactly 6 e- domains has sp 3 d 2 hybrids.

34. Non-bonding e- pairs Lone pairs occupy hybrid orbitals, too Ex: H 2 O (g): observed as bent; but e- domain is tetrahedral 2p 2s Four sp 3 orbitals (2 bonding, 2 non-bonding) O (ground) → O (sp 3 hybrid) 2 non-bonding pairs (lone pairs) 2 bonding pairs H

35. 9.6: Multiple Bonds Draw Lewis structures. For C’s: label hybridization, molecular geometry, and unique bond angles C 2 H 6 C 2 H 4 C 2 H 2 C 6 H 6

36. Sigma (  ) bonds in C 2 H 4 Ex: ethene; C-C  -bonds and C-H  -bonds result from axial overlap of H s-orbitals and C sp 2 - orbitals

37. Pi (  ) bonds in C 2 H 4 Each C has 4 valence e - : – 3 e - for 3  bonds – 1 e - for 1  bond, which results from side-by-side overlap of one non-hybridized p-orbital from each C 2p C 2s sp 2 hybrids bond axially =  bonds p orbital bonds side-by-side =  bond

38. Sigma (  ) bonds in C 2 H 2 Ex: ethyne (a.k.a. acetylene) C-C  -bond and C-H  - bonds result from axial overlap of H s-orbitals and C sp-orbital

39. Pi (  ) bonds in C 2 H 2  Each C has 4 valence e - : – 2 e - for 2  bonds – 2 e - for 2  bonds, which result from side-by-side overlap of two non-hybridized p-orbitals from each carbon sp hybrids bond axially =  bonds 2p C 2s p orbital bonds side-by-side =  bonds 

40. Sigma (  ) bonds in C 6 H 6 Ex: benzene; C-C  -bonds and C-H  -bonds result from axial overlap of H s-orbitals and C sp 2 - orbitals http://www2.chemistry.msu.edu/faculty/reusch/VirtTxtJml/intro3.htm#strc8c

41. Localized v. Delocalized Bonds Delocalized bonds are present in compounds showing resonance structures, while electrons are localized in most other bonds.

42. Localized vs. Delocalized  Bonds (localized)(delocalized – MINIMUM OF 4 c’S)

43. Delocalized  bonds in C 6 H 6 C-C  -bonds result from overlap of one non-hybridized p-orbitals from each C Delocalization of e - in  -bonds results in a “double-donut” shaped e - cloud above and below the molecular carbon plane.

44. Limitations of Hybrid Orbital Theory Hybrid orbital theory assumes that all bonds are formed with localized electrons, which is not true. MO (Molecular Orbital) theory explains bonding in terms of delocalized orbitals as well.

45. 9.7: Molecular Orbital (MO) theory  So far we have used valence-bond theory (covalent bonds form from overlapping orbitals between atoms) with hybrid orbital theory and VSEPR theory to connect Lewis structures to observed molecular geometries. However, VB theory does not explain the magnetic or spectral properties of a molecule.  MO theory is similar to atomic orbital (AO) theory (s, p, d, f orbitals) and helps to further explain some observed phenomena, like unpredicted magnetic properties in molecules like those in O 2.  AO are associated with the individual atoms, but MO are associated with the whole molecule.

46. *AO & MO in H 2  Combination of two 1s AO from each H forms two MO in H 2 molecule.  Bonding MO: form between nuclei and are stable  Antibonding MO: marked with *; form “behind” nuclei and are less stable. 1s Molecular orbitals E  1s  * 1s Atomic orbitals Bonding orbital Anti-bonding orbital

47. *Types of MO Sigma (  ) MO: form from combinations of: – Two 1s or 2s orbitals from different atoms; written as  1s or  2s. – Two 2p z orbitals from different atoms (axial overlap); written as  2p z.(Some sources say 2 px orbitals?) Pi (  ) MO: form from combinations of: –Two 2p x or 2p y orbitals from different atoms; written as  2p x or  2p y. –Do not appear until B 2 molecule

48. MO s from Atomic p-Orbital Combinations  P orbitals can interact with each other forming either sigma molecular orbitals,  2p, in a end-to-end overlap or pi molecular obrbitals,   p, in a side-to-side overlap.  The order of energy for MO s derived from 2p orbitals is  2p <   p <   p* <  2p*  There are three perpendicular p orbitals, so two sigma p orbitals (one bonding and one antibonding) and four pi p orbitals (two bonding and two antibonding) are formed.  This energy order gives the expected MO diagram for most of the p-block elements for homonuclear diatomic molecules.

49. MO s for B, C and N  The energy order of p orbitals results from the assumption that since s and p orbitals have differences in energy, they do not interact with each other. (or mix)  However, when 2p atomic orbitals are half filled, such as in B, C and N, the repulsions between e are little and the energy of these p orbitals is not much different than the s atomic orbital, which leads to s and p orbital mixing. This mixing lowers the energy of the 2s bonding and antibonding orbitals and increases the energy of sigma 2p (bonding and antibonding) orbitals.The pi 2p orbitals are not affected. This mixing gives a different energy order:   2s <  2s* <   p <  2p <   p* <  2p

50. *MO diagrams for “< O 2 ” Resulting MO for diatomic molecules with < 16 e- (B 2, C 2, N 2, etc.) Bond order = ½ (# bonding e - - # antibonding e - ) B.O. (N 2 ) = ½ (10 – 4) = 6 / 2 = 3 (triple bond) N 2 has no unpaired electrons which makes it diamagnetic. N atom

51. *MO diagrams for “≥ O 2 ”  Resulting MO for diatomic molecules with ≥ 16 e- (like O 2, F 2, Ne 2, etc.) O atom Bond order = ½ (# bonding e - - # antibonding e - ) B.O. (O 2 ) = ½ (10 – 6) = = 2 (double bond) O 2 has unpaired electrons which makes it paramagnetic.

52. Liquid N 2 and liquid O 2 From U. Illinois: http://www.chem.uiuc.edu/clcwebsite/liquido2.html N2N2 O2O2

53. Magnetism In an element or compound:  Diamagnetism: all e - paired; no magnetic properties  Paramagnetism: at least 1 unpaired e - – * Drawn into exterior magnetic field since spins of atoms become aligned; unlikely to retain alignment when field is removed Ex:NO ScMn O 2  * Ferromagnetism: occurs primarily in Fe, Co, Ni – Drawn into exterior magnetic field since spins of atoms become aligned; very likely to retain alignment when field is removed (i.e., “a permanent magnet”) – Nd 2 Fe 14 B is very ferromagnetic