CHAPTER 7 Molecular Geometry, Intermolecular Forces, and Bonding Theories

1. CHAPTER 7 Molecular Geometry, Intermolecular Forces, and Bonding Theories

2. Shapes of Molecules and Ions Molecules and ions have a three dimensional shape. This shape can be important in determining the chemical behavior of a substance. We will discuss two types of geometries: 1) electron domain geometry - The arrangement of electron containing regions about a central atom. This is sometimes called electron cloud geometry. 2) molecular geometry - The arrangement of atoms about a central atom. The number of electron containing regions about a central atom is equal to the number of lone pair regions + the number of covalent bonds. Note that in counting electron containing regions a single, double, or triple bond counts as one region.

3. Examples of Counting Electron Containing Regions How many electron containing regions are there for central atoms in each of the mole-cules at left?

4. VSEPR Theory We may predict both electron cloud geometry and molecular geometry using VSEPR (Valence Shell - Electron Pair Repulsion) theory. The idea behind this theory is that electron containing regions will arrange themselves about a central atom in such a way as to put those regions as far apart as possible. This is due to the repulsive force existing between particles of the same charge (in this case electrons). The various common cases for electron and molecular geometry are given in Table 7.2.

5. Two Regions If we have two electron containing regions around a central atom, then placing them opposite one another keeps them as far apart as possible. In this case, both the electron geometry and the molecular geometry are linear.

6. Three Regions For three regions about a central atom the electron geometry is trigonal planar (120 angle between the regions). There are two possible molecular geometries: If all three regions are bonds - trigonal planar If two of the three regions are bonds, and one is a lone pair of electrons - nonlinear (bent)

7. Four Regions For four regions about a central atom the electron geometry is tetrahedral (109 angle between the regions). There are three possible molecular geometries: 4 bonds - tetrahedral 2 bonds - nonlinear (bent) 3 bonds - trigonal pyramid

8. Five Regions For five regions about a central atom the electron geometry is trigonal bipyramid. In this case not all of the angles between electron containing regions are the same. Instead, we may divide the regions into equitorial regions and axial regions.

9. Five Regions There are four possible molecular geometries: 5 bonds - trigonal bipyramid 3 bonds - T-shape 4 bonds - “see-saw” 2 bonds - linear

10. Six Regions For six regions about a central atom the electron geometry is octahedral. The bond angles between adjacent regions are all 90.

11. Six Regions There are three common molecular geometries: 6 bonds - octahedral 5 bonds - square pyramid 4 bonds - square planar

12. Example: What are the electron geometries and molecular geometries around the central atom in BrF3, POCl3, and H2SO3?

13. Summary electron cloud bonds electron molecular regions geometry geometry 2 2 linear linear 3 3 trigonal planar trigonal planar 2 nonlinear (bent) 4 4 tetrahedral tetrahedral 3 trigonal pyramid 2 nonlinear (bent)

14. Summary electron cloud bonds electron molecular regions geometry geometry 5 5 trig. bipyramid trig. bipyramid 4 see-saw 3 T-shape 2 linear 6 6 octahedral octahedral 5 square pyramid 4 square planar

15. Geometry in Large Molecules For large molecules we can discuss the electron cloud and molecular geometry around each one of the interior atoms. Example: Give the electron cloud and molecular geometry for each interior atom in the molecule glycine (NH2CH2COOH).

16. Deviations From “Pure” Geometries We have assumed that the angles observed in molecular geometries do not depend on how many bonds and lone pair regions there are. In fact, the number of regions containing lone pair electrons has a small effect on the observed bond angles. In general we may say the repulsive forces between electron containing regions in a molecule or ion are stronger for lone pair electrons than for bonding electrons. So lone pair - lone pair > lone pair - bonding > bonding - bonding This is due to the fact that bonding electrons are more localized than lone pair electrons, as they must appear between bonded atoms.

17. In the molecules at left (all of which have four electron containing regions around a central atom) the bond angle decreases from 109.5  in CH4 (pure tetrahedral geometry) to 107.0  in NH3 to 104.5  in H2O. In CH2O (below) all of the bond angles would be 120. ° in pure trigonal planar geometry.

18. Representing Three Dimensional Structure Molecules have a three dimensional structure. We need a systematic method for representing these structures in two dimensions. We do this as follows For convenience, we will sometimes used a dashed line to replace the hatched wedge for a bond going into the page. Example: Give the three dimensional structures for CH3Cl and PF5.

19. Polar Molecule A polar molecule is a molecule where the center of positive and negative charge do not coincide. Two things are required for a molecule to be polar 1) The molecule must have at least one polar bond. 2) The contributions from the polar bonds cannot cancel.

20. For a polar covalent bond the difference in electronegativity between the bonded atoms should be about 0.5 or greater. This means that carbon-carbon bonds are completely nonpolar, and carbon-hydrogen bonds are only slightly polar. EN(C) = 2.5 EN(H) = 2.1 The table of electronegativities and the molecular geometry for the molecule can be used to determine whether the molecule is polar or nonpolar. The polarity of a molecule is expressed in terms of the dipole moment of the molecule. The symbol for dipole moment is . Dipole moments are measured in units of Debye (D).

21. Example Consider the following molecules: H2O, CF4, NH3, and CH3COCH3. Which of these molecules is expected to have a permanent dipole moment?

22. Intermolecular Forces Intermolecular forces are the forces that exists between different molecules or particles. We are more concerned with long range attractive forces and will ignore short range repulsive forces. Ion-ion - The attractive force acting between cations and anions. These are strong, and are found in substances where ionic bonding occurs..

23. Dipole-Dipole Forces Dipole-dipole - The attractive force acting between polar molecules. The attraction is between the partial positive charge (+) on one molecule and the partial negative charge (-) on a different molecule. Generally speaking, the larger the partial positive and negative charges the stronger the dipole-dipole attraction.

24. Dipole-Dipole Forces and Boiling Point When molecules have strong intermolecular attractive forces it takes more energy to overcome those attractive forces. One way of seeing this is in the boiling point for a substance. Generally speaking, the stronger the dipole-dipole attraction between molecules the higher the boiling point, parti-cularly for substances with approximately the same molecular mass.

25. Hydrogen Bonding Hydrogen bonding - A particularly strong form of dipole-dipole attractive force. It is the attractive force that exists between a hydrogen atom bonded to an N, O, or F atom and lone pair electrons on a different N, O, or F atom (often an atom in a different molecule).

26. Evidence For Hydrogen Bonding One effect of hydrogen bonding is to raise the boiling point of a liquid. This occurs because it requires more energy (and so a higher temperature) to break apart strong attractive forces between molecules than it does to break apart weak attractive forces between molecules. substanceboiling pointhydrogen bonding? H2O 100.0 C yes H2S - 60.7 C no H2Se - 41.5 C no H2Te - 4.4 C no

27. Boiling Points for Binary Hydrogen Compounds

28. London Dispersion Forces London dispersion forces - The attractive force that is due to the formation of instantaneous dipoles in a molecule. These instantaneous dipoles arise from the random motion of the electrons in the molecule.

29. Strength of London Dispersion Forces London dispersion forces are present in all molecules, but are the only intermolecular force present in nonpolar molecules. The strength of London dispersion forces is approximately proportional to the number of electrons, and so to the size of the molecule. Therefore, as a general rule, the larger the molecule the stronger the London dispersion forces. substanceboiling point He - 268.6 C Ne - 245.9 C Ar - 185.7 C Kr - 152.3 C Xe - 107.1 C

30. Induced Dipole and Polarizability London dispersion forces are closely related to a second property of molecules, called polarizability. The polarizability of a molecule refers to the extent to which the electron cloud distribution is distorted when a positive or negative charge is brought close to the molecule. In general, large molecules, and molecules containing atoms whose valence electrons are far away from the atomic nucleus are more polarizable than other molecules. The more polarizable a molecule the stronger the London dispersion forces it experiences.

31. Ion-Dipole Forces Ion-dipole - The attractive force between an ion and a polar molecule. Responsible for the dissolution of some ionic substances in polar liquids such as water. In general, the smaller the ion and the larger the charge the stronger the ion-dipole attractive force. Solvation - The close association of solvent molecules with solute molecules or ions. Hydration - Solvation when the solvent is water.

32. Summary of the Types of Intermolecular Forces Ion – ion. Forces between cations and anions. Dipole – dipole. Forces between molecules with a permanent dipole moment. This category includes hydrogen bonding, a particularly strong type of dipole – dipole force. London dispersion forces. Due to random movement of electrons. All particles have this type of force, but it is most important in molecules with no permanent dipole moment. Mixed forces: Ion – dipole is the most important, but this also includes ion – induced dipole and dipole – induced dipole. It is responsible for the solubility of some ionic compounds in water. We will discuss these and related forces in detail when discussing solution formation.

33. Theories For Molecular Bonding The Lewis dot structure method predates quantum mechanics. While it is a good qualitative description of covalent bonding it fails to work well in some cases (resonance structures are one example). There are two general methods that use quantum mechanics to improve on the Lewis dot structure method: Valence Bond theory - Discusses bond formation in terms of overlap of atomic orbitals (often hybrid orbitals). Molecular Orbital theory - Uses quantum mechanics to solve the Schrodinger equation for a molecule or ion. This is a very mathematical approach.

34. Valence Bond Theory For H2 In valence bond theory bond formation is pictured as occurring due to the overlap of atomic orbitals of the atoms that are bonded together. Each of the orbitals contains one electron. When the orbitals overlap, the electron from each atomic orbital is shared by the two atoms. Notice the electrons need to be opposite spin. The covalent bond forms from the overlap of the two 1s atomic orbitals of the hydrogen atoms that are bonded together. This places the electrons between the two hydrogen nuclei, which holds the molecule together.

35. Bond Formation and Energy The formation of a covalent bond between two atoms occurs because it leads to a lower energy than exists for separate atoms. This can be seen in the bonding of two H atoms to form an H2 molecule.

36. Shortcoming of Simple Valence Bond Theory The above picture breaks down when you need to form several different valence bonds for the same atom out of orbitals of different types. Example: Be in BeH2 Be 1s2 2s2 H 1s1 We want the beryllium atom to contribute one electron to each bond between Be and an H atom, but the 2s orbital of beryllium is full, with no unpaired electron. So what do we do?

37. Hybrid Orbitals In this case, before we form the valence bonds we first construct a new set of equivalent orbitals, called hybrid orbitals (in this case sp hybrid orbitals). We can envision the process taking place as indicated below. While the process of forming hybrid orbitals raises the energy of the electrons in Be, this is more than made up for when the two Be - H bonds are formed.

38. Formation of sp Hybrid Orbitals The sp hybrid orbitals are formed from combinations of the s and px atomic orbitals. Since we begin with two atomic orbitals, we end up with two hybrid orbitals.

39. Use of sp Hybrid Orbitals For Bonding (BeH2) The two sp hybrid orbitals formed in Be can now be used to form the two Be - H valence bonds.

40. sp3 Hybrid Orbitals In methane (CH4) we need the central carbon atom to have four equivalent hybrid orbitals to form the four C - H bonds. This is done using sp3 hybridization. CH4. C: 1s2 2s2 2p2 As before, we start by “unpairing” the electrons in the carbon atom by promoting one electron from a 2s to a 2p orbital __ __ __ __ promote __ __ __ __ 2s 2p 2s 2p

41. sp3 Hybrid Orbitals (Continued) Now that we have one electron in our 2s atomic orbital, and in each of the three 2p atomic orbitals, we combine the four orbitals to form our set of four sp3 hybrid orbitals. Since we began with four atomic orbitals, we end up with four hybrid orbitals.

42. These sp3 hybrid orbitals can now be used to form valence bonds or to hold lone pairs of electrons.

43. Naming and Use of Hybrid Orbitals We name hybrid orbitals by listing the type of atomic orbitals used to construct the hybrid orbitals (s, p, or d). If we use more than one of a particular type of orbital, we indicate the number of orbitals used by a superscript to the right of the symbol for the orbital. There is a simple relationship between the number of electron containing regions around a central atom and the type of hybrid orbitals that are required. Number of regionsHybrid orbitalsElectron geometry 2 sp linear 3 sp2 trigonal planar 4 sp3 tetrahedral 5* sp3d trigonal bipyramid 6* sp3d2 octahedral * Not possible for atoms from 2nd row of periodic table

44. Why PF5 Exists and NF5 Does Not We can now return to an observation we made in the last chapter, that in looking at molecules made from a group 5 nonmetal and fluorine the molecule PF5 exists, but NF5 does not. We can see why this is the case by looking at the hybridization required for the central atom in these two molecules. N [He] 2s2 2p3 P [Ne] 3s2 3p3 3d0 __ __ __ __ __ __ __ __ __ promote __ __ __ __ __ __ __ __ __ 3s 3p 3d Hybridize to get 5 sp3d hybrid orbitals. This is not possible for nitrogen because there are no 2d orbitals, and so hybrid orbitals requiring use of one or more d orbitals cannot be formed.

45. Appearance of Hybrid Orbitals Each particular type of hybrid orbital has its own geometry, which in all cases corresponds to the geometries predicted using VSEPR theory. Also notice that in all cases the number of atomic orbitals we begin with is equal to the number of hybrid orbitals we end up with.

46. Sigma () and Pi () Bonds The above description works well for single bonds between atoms. However, when we have a multiple bond (double or triple bond), the bonds can be divided into two types: sigma bond - Formed from the overlap of atomic or hybrid orbitals; places electrons directly between the bonded atoms. pi bond - Formed from the overlap of p type atomic orbitals; places electrons above and below the region between the bonded atoms. In general, a single bond will always be a sigma bond. For a multiple bond, one of the bonds will be a sigma bond and the other bonds will be pi bonds.

47. Example: C2H4 (ethene) (C atoms are sp2 hybridization)

48. Summary Total number Number of Number of bonds sigma bonds pi bonds 1 1 0 2 1 1 3 1 2 So the first bond is always a sigma bond, while any additional bonds are pi bonds.

49. Example How many sigma bonds and how many pi bonds are present in the molecule shown below?

50. Pi Bonds and Free Rotation Because pi bonds are formed from the overlap of p orbitals, molecules that contain pi bonds (molecules with double or triple bonds) cannot rotate freely around those bonds, since to do so would mean breaking one or more pi bond. In ethane (C2H6) rotation around the C - C bond does not affect the bond and so is allowed. In ethene (C2H4) rotation around the C = C bond breaks the pi bond, and so will not occur (except at high energy).