1. Base Pairing in DNA

2. Red = O Grey = C White = H Purple = K Ionic Radii Li + = 0.68 Å Na + = 0.97 Å K + = 1.33 Å Rb + = 1.47 Å Cavity Size (O-O Dist.) = 1.40 Å K + fits best Crown Ether, C 12 H 24 O 6 (18-Crown-6)

12. Rules for Predicting Molecular Geometry 1. Sketch the Lewis structure of the molecule or ion 2. Count the electron pairs and arrange them in the way that minimizes electron-pair repulsion. 3. Determine the position of the atoms from the way the electron pairs are shared. 4. Determine the name of the molecular structure from the position of the atoms. 5. Double or triple bonds are counted as one bonding pair when predicting geometry.

14. Note: The same rules apply for molecules that contain more than one central atom

15. The Dipole A dipole arises when two electrical charges of equal magnitude but opposite sign are separated by distance.

16. The dipole moment (  )  = Qr where Q is the magnitude of the charges and r is the distance

17. The sum of these vectors will give us the dipole for the molecule For a polyatomic molecule we treat the dipoles as 3D vectors

18. Overlap of Orbitals

19. The point at which the potential energy is a minimum is called the equilibrium bond distance The degree of overlap is determined by the system’s potential energy equilibrium bond distance

20. 2s These new orbitals are called hybrid orbitals The process is called hybridization What this means is that both the s and one p orbital are involved in bonding to the connecting atoms Formation of sp hybrid orbitals The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals.

21. Formation of sp 2 hybrid orbitals

22. Formation of sp 3 hybrid orbitals

23. Hybrid orbitals can be used to explain bonding and molecular geometry

24. Multiple Bonds Everything we have talked about so far has only dealt with what we call sigma bonds Sigma bondA bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms. Sigma bond (  )  A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.

25. Pi bondA bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis). Pi bond (  )  A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

26. Example: H 2 C=CH 2

28. Example: HC  CH

29. Delocalized  bonds When a molecule has two or more resonance structures, the pi electrons can be delocalized over all the atoms that have pi bond overlap.

30. In general delocalized  bonding is present in all molecules where we can draw resonance structures with the multiple bonds located in different places. Benzene is an excellent example. For benzene the  orbitals all overlap leading to a very delocalized electron system Example: C 6 H 6 benzene

31. The one that is lower in energy is called the bonding orbital, The one higher in energy is called an antibonding orbital. These two new orbitals have different energies. BONDING ANTBONDING Moleculuar Orbital (MO) Theory

32. Energy level diagrams / molecular orbital diagrams

33. MO Theory for 2nd row diatomic molecules Molecular Orbitals (MO’s) from Atomic Orbitals (AO’s) 1. # of Molecular Orbitals = # of Atomic Orbitals 2. The number of electrons occupying the Molecular orbitals is equal to the sum of the valence electrons on the constituent atoms. 3. When filling MO’s the Pauli Exclusion Principle Applies (2 electrons per Molecular Orbital) 4. For degenerate MO’s, Hund's rule applies. 5. AO’s of similar energy combine more readily than ones of different energy 6. The more overlap between AOs the lower the energy of the bonding orbital they create and the higher the energy of the antibonding orbital.

34. Example: Li 2

35. MOs from 2p atomic orbitals 1) 1 sigma bond through overlap of orbitals along the internuclear axis. 2) 2 pi bonds through overlap of orbitals above and below (or to the sides) of the internuclear axis.  

37. Interactions between the 2s and 2p orbitals The  2s and  2p molecular orbitals interact with each other so as to lower the energy of the  2s MO and raise the energy of the  2p MO.

38. For B 2, C 2, and N 2 the interaction is so strong that the  2p is pushed higher in energy than  2p orbitals

39. Paramagnetism of O 2