1. Molecular Geometry & Bonding Theories Chapter 9

2. Molecular Shapes  Lewis Structures that we learned do not tell us about shapes, they only tell us how many bonds and what type of bonds.  The actual shape is determined by a molecules bond angles  The bond angles, along with bond lengths determine the shape and size of the molecule.

3. VSEPR Model  The shape of a molecule can greatly affect its properties.  Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

4. Balloons  1 balloon  2 balloons  3 balloons  4 balloons  Electron domains- where the electrons are most likely to be found  The best arrangement of a given number of electrons is the one that minimizes the repulsions among them.

5. VSEPR  Molecules take a shape that puts electron pairs as far away from each other as possible.  Have to draw the Lewis structure to determine electron pairs.  bonding  nonbonding lone pair  Lone pair take more space.  Multiple bonds count as one pair.

6. VSEPR  The number of pairs determines  bond angles  underlying structure  The number of atoms determines  actual shape

7. VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal

8. Molecule Shape  The shape of a molecule can be related to these 5 basic arrangements of electron domains  The arrangement of electron domains around its central atom is called its electron domain geometry  Use electron domain geometry to predict molecular geometry (NH 3 )

9. How to do it  1. Sketch the Lewis structure of the molecule or ion  2. Count the total number of electron domains around the central atom and arrange them in a way that minimizes electron repulsion (greatest bond angles)  3. Describe the molecular geometry in terms of the angular arrangement of the bonded atoms  4.Note that a double or triple bond counts as only one electron domain (CO 2 )

10. Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

11. Actual Shape Electron Bonding Non Pairs Pairs Bonding Pairs Shape Pairs Shape 5 5 0 Trigonal bipyramidal bipyramidal 5 4 1 Seesaw 5 3 2 T-shaped 5 2 3 Linear

12. Actual Shape Electron Bonding Non Shape Pairs Pairs Bonding Pairs Pairs 6 6 0 Octahedral 6 5 1 Square pyramidal pyramidal 6 4 2 square planar planar 6 3 3 T-shaped 6 2 4 Linear

13. Want to try some?  SF 4  O 3  CH 4  CH 3 COOH  SnCl 3 -  HBr

14. Some reminders  Non bonding electron pairs have a greater repulsive force, and so may cause small deviations in bond angle  Multiple bonds also have a greater repulsive force due to a higher concentration of electronic charge, and so they will also cause slight deviations in bond angles  Also when calculating the number of electron domains, it is always with respect to the central atom.

15.  For molecules with 5 electron domains, any lone pairs will first occupy the equatorial positions because they are at the furthest angle from all other electron domains

16. No central atom  Can predict the geometry of each angle.  build it piece by piece.  Think of the molecule as having several different central atoms and calculate each of the geometries individually  Consider acetic acid (CH 3 COOH)

17. Polarity of polyatomic molecules  Remember?  Bond Polarity- measure of how equally the electrons are shared in a bond  Dipole moment- a quantitative measure of the amount of charge separation in a molecule  This was in terms of each single bond, how does this carry over when talking about all the bonds in a molecule?

18.  For each bond in a molecule you can consider the bond dipole which is dependent on only the two atoms in the bond.  Consider CO 2, the C=O bond is polar, because each C=O bond is the same, they are both considered polar. This does not mean that the whole molecule is polar, just the individual bonds.  Why Not?

19.  Bond dipoles are vector quantities (have a magnitude and direction)  The dipole of a polyatomic molecule is the sum of all the component dipoles  If you think about the molecular geometry (shape) of CO 2, even though the polarity of the C-O bonds are equal in magnitude, they are opposite in direction, and so the whole CO 2 molecule will have a 0 dipole  A CO 2 molecule is said to be non polar, even though it has polar bonds!

20. So, we must consider the geometry!  Consider H 2 O  Has polar bonds, but since the molecule is bent, these bonds do not cancel out, and so we can say that water is a polar molecule.

21. Try to make it simple Remember!!!!!! Bond type does not translate to molecule type. Bond Symmetry Molecule Type Type Non Polar both Non Polar Polar non- symmetrical Polar Polar symmetrical Non Polar

22. Covalent Bonding & Orbital Overlap  Valence Bond Theory- tries to explain why and how bonds occur between atoms.  Covalent Bonding occurs when atoms share electrons. There is a concentration of electronic charge that builds up between the two nuclei  The Lewis theory says this occurs when the valence atomic orbitals of one atom merge with those of another (they overlap)

23. Covalent Bonding continued  H 2 Each Hydrogen has a single electron in a 1s orbital. When the orbitals overlap, the two electrons are concentrated between both nuclei and attracted to both, holding the atoms together. This is how a covalent bond is formed.  HCL- same idea. The Hydrogen has 1 unpaired electron in the 1s shell, and the chlorine has one unpaired electron in the 3p shell. They overlap and the electrons occupy a common space.

24.  There is always an optimal distance between two bonded nuceli in any covalent bond.  This is what dictates bond length- when the distance at which the attractive forces between the unlike charges are balanced by the repulsive forces of the like charges

25. Polyatomic molecules and orbital overlap  To apply the valence bond theory, you must consider the formation of electron pairs and the shape of the molecules

26. Hybrid orbitals  BeF 2  Important to remember that both bonds should be the same. So, how can we explain the bonding in terms of orbital overlap.  Be has electron configuration 1s 2 2s 2 so no unpaired electrons. Each Fluorine has 1 unpaired electron in the p orbital.

27.  First need to promote an electron of Be (excite- requires energy)  Now you have two unpaired electrons. This is great! These two electrons are available to interact with the two unpaired electrons available from both the fluorines.  There is a problem with this explanation of the BeF 2 model!

28.  Both the bonds should be the same. If one of the bonds involves an s orbital overlapping with a p orbital, and the other one involves a p orbital overlapping with another p orbital, we would not guess that these are the same bonds.  Something called hybridization occurs. Instead of having 1 electron left in an s, and 1 in a p, the two orbitals come together and form a new orbital called an sp hybrid orbital. The process is called hybridization.

29. sp, sp 2, sp 3  If you mix a 2s orbital with a 2p orbital, you will get (2) sp hybrid orbitals  If you mix a 2s orbital with 2 2p orbitals this will make (3) sp 2 hybrid orbitals (all equivalent)  If you mix a 2s orbital with (3) 2p orbitals, you will make (4) equivalent sp 3 orbitals.

30. Orbitals- Shape  So far we have used the idea of orbital overlap to try and explain why atoms bond together the way they do, for example why does carbon typically bond to four hydrogen atoms  If we look at the actual shape of the hybrid orbitals, we will see that the type of hybridization is going to depict the type of molecular geometry the molecule will exhibit.

31. Hybridization and the d orbital  Only occurs after period 2 because of the availability of the d shell to form an expanded valence shell  (5) sp 3 d - trigonal bipyramidal  (6) sp 3 d 2 - octahedral

32. Not definitive, only a prediction  Use the following steps to predict the hybrid orbitals used by an atom in bonding  1. Draw the Lewis Structure for the molecule or ion  2. Determine the electron-domain geometry using the VSEPR model  3. Specify the hybrid orbitals needed to accommodate the electron pairs based on their geometric arrangement

33. Multiple Bonds  Internuclear Axis-line between the nuclei  Sigma bonds seen in most single bonds. The concentration of electrons is directly between the involved atoms, lying on the internuclear axis  In multiple bonds 2 p orbitals are oriented perpendicularly to the internuclear axis. The resulting bond is a π bond.

34. Multiple Bonds continued  The basic difference between a π bond and a sigma bond is that in a π bond the electron density lies above and below the internuclear axis.  Π bonds are generally weaker than sigma bonds because there is less of an overlap of the two domains  A double bond usually consists of one pi bond and one sigma bond  A triple bond usually consists of one sigma bond and two pi bonds.

35. C2H4C2H4C2H4C2H4  Three electron domains, or bond angles of 120° would lead us to believe this carbon underwent sp 2 hybridization.  If this were true than each carbon would have one electron unhybridized in the 2p shell.  The interaction between these two unpaired p would result in a π bond lying above the plane.

36. C2H2C2H2C2H2C2H2  The bond angles of 180°, would lead us to believe that each of these carbons underwent sp hybridization. This would leave each carbon with 2 unpaired p orbitals. Now, you can see how the triple bond is made up of two different types of bond. One sigma bond coming from the sp hybridization, and two pie bonds coming from the unpaired p orbitals.

37. Localized or delocalized  Molecules that exhibit resonance structures are said to have delocalized π bonds  In localized bonds, the electrons involved in bonding come directly from the 2 atoms involved  In delocalized bonds, the bonding is affected by neighboring atoms

38. Benzene  The best way to describe benzene and its π bonds is to say that the electrons involved in benzene are smeared out across the entire atom.  Accounts for the extremely stable configuration of benzene

39. Back to basics  All bonded atoms have at least one sigma (localized) bond  Sigma bonds are localized in the regions between the atoms, and do not contribute to the bonding between other atoms  When atoms share more than one pair of electrons, the additional bonds are π bonds that do not lye in the region.  Molecules with resonance structures often have their bonds delocalized over all the atoms

40. Molecular Orbitals  Molecular Orbital Theory- another  Molecular Orbital theory  MO- associated with the entire molecule, based on the wave theory

41.  Solve the equations for H 2  H A H B  MO 1 = 1s A - 1s B  MO 2 = 1s A + 1s B

42. MO  Sigma MO- electron density lies on the internuclear axis  * indicates antibonding  Energy level diagrams  Each MO can hold two electrons of opposite spin (Pauli Exclusion Principle)  Molecular Orbital Theory accounts for why some atoms occur as diatomic molecules, and some do not

43. Bond Order  The stability of a covalent bond is related to its bond order.  Bond order (# bonding e - - # non-bonding e - ) 2 A bond order of 0 means that no bonds exist

44. Second Row Diatomics and MO’s  # MO’s formed is the same as # atomic orbitals combined  Only similar energy orbitals will combine  Look at Li 2  Bond order says it has a single bond  The bond comes mostly from the 2s electrons, accounting for an important rule we learned early on- core electrons usually do not contribute to bonding

45. Don’t think about 1s  If you are doing MO’s for a second row element, don’t consider the 1s as we have already established, it does not contribute to bonding  Be 2 - does not exist, has a bond order of 0  Now, what about the 2p?

46.  Sigma and pi MO’s for the same reason as atomic orbitals  Degenerate orbitals- molecular orbitals that have the same energy (two π 2p )

47.  Energy Level diagrams have assumed the 2s and 2p orbitals do not interact with each other