1. Molecular Geometry and Bonding Theories

2. Physical and chemical properties of a molecule are determined by: size and shape strength and polarity of bonds Lewis structures do not indicate shapes of molecules, simply the number and types of bonds. To translate Lewis structures into three dimensions, bond angles- the angles made by lines joining the nuclei of the atoms in the molecule- must be used.

4. Valence Shell Electron Pair Repulsion Theory (VSEPR) Best arrangement of electron pairs is the one that minimizes repulsions. Arrangement of electron pairs around a central atom is called electron-pair geometry. No distinction is made between bonding and nonbonding electrons. Molecular geometry is the arrangement of atoms in space; distinction between bonding and nonbonding electrons.

5. How do electron pairs affect shape? 2 electron pairs = linear = 180 o 3 e.p. = trigonal planar = 120 o 4 e.p. = tetrahedral = 109.5 o 5 e.p. = trigonal bipyramidal = 120 o and 90 o

6. 6 e.p. = octahedral = 90 o In order to determine electron pair geometry, look at number of electron pairs attached to central atom and don’t distinguish between bonding electrons and lone pairs. Molecular geometry takes into account lone pairs which influence the shape of the molecule. Multiple bonds are treated the same as single bonds.

9. Bonding Non bonding GeometryPairs 2 0 Linear 3 0 Trigonal Planar 2 1 Angular 4 0 Tetrahedral 3 1 Trigonal Pyramidal 2 2 Angular

10. 50Trigonal Bipyramidal 41Seesaw 32T-shaped 23Linear 60Octahedral 51Square Pyramidal 42Square Planar

11. Determine the electron pair geometry and molecular geometry for the following: SF 4 IF 5 ClF 3 CO 3 2- H 2 S

12. http://www.dcu.ie/~pratta/jmgallery/JGALLERY.HTM

13. Non-bonding electrons exert greater repulsive forces on adjacent pairs and compress angles. CH 4, NH 3, H 2 O all have a tetrahedral electron pair geometry, but their molecular geometry is dictated by the presence of lone pairs on the central atom. Bond angles are 109.5 o, 107 o, and 104.5 o respectively. In absence of central atom, predict geometry around each atom of backbone.

15. Polar molecules -degree of polarity measured by dipole moment -dipole moment of molecule depends on the polarities of the individual bonds and the geometry of the molecule. Once you have established whether the individual bonds are polar, look at the symmetry of the molecule, if symmetric, non-polar, if asymmetric, polar. Lone pairs on the central atom = polar.

18. VSEPR Theory Provides simple means for predicting shapes of molecules. Does not explain why bonds exist or form. http://www.chem.purdue.edu/gchelp/vsep r/http://www.chem.purdue.edu/gchelp/vsep r/ Molecular shapes

19. Valence Bond Theory Combines Lewis’s idea of electron pair bonds with atomic orbitals. Atomic orbital of one atom merges with that of another atom Orbitals share a region of space or overlap.

20. http://www.mhhe.com/physsci/chemistry/essential chemistry/flash/hybrv18.swf Hybridization tutorial

21. How does valence bond theory explain molecules like BeF 2 ? Be = 1s 2 2s 2 F = 1s 2 2s 2 2p 5 In order for Be to bond with 2 F atoms, hybridization or mixing of the orbitals occurs. Hybrid orbitals require energy but they can overlap more strongly resulting in a stronger bond. The energy released offsets the energy expended in the formation of the hybrid orbital.

24. Link electron pair geometry with hybridization: linearsp trigonal planarsp 2 tetrahedralsp 3 trigonal bipyramidalsp 3 d octahedral sp 3 d 2 Predict the hybridization of the following: NH 2 - SF 4 SO 3 2- SF 6

25.  bonds are formed by the overlap of two s orbitals. Concentration is symmetric on internuclear axis.  bonds can form by the overlap of two s orbitals, an s and a p orbital or two p orbitals that are facing each other. All single bonds are  bonds.

26.  bonds are formed by the side-ways overlap of p orbitals.  bonds are oriented perpendicular to the internuclear axis. Because there is less overlap in a  bonds, generally these bonds are weaker than  bonds. When multiple bonds are formed (such as double and triple bonds), the first bond is a  bond and the remaining bonds are  bonds. Predict the hybridization and the number of  and  bonds in formaldehyde: H 2 CO.

28.  and  bonds are considered localized: electrons are associated with the two atoms forming the bond. A molecule that does not have localized electrons is benzene, C 6 H 6. Benzene has two resonance structures resulting in 6 bonds of equal length. The 3  bonds that form are said to be delocalized among the 6 carbon atoms. It is important to understand that wherever resonance occurs with multiple bonds, the  bonds that form will be considered delocalized.

30. General Conclusions Every pair of bonded atoms shares one or more pairs of electrons. In every bond at least one pair of electrons is localized. The electrons in a  bonds are localized. When atoms share more than one pair of electrons, the additional pairs form  bonds. Electrons in  bonds that extend over more than two atoms are delocalized.

31. Molecular Orbital (MO) Theory Molecular orbitals form from a combination of atomic orbitals. Contain a maximum of two electrons. Two atomic orbitals form two molecular orbitals.

34. When 2 hydrogen atoms combine, the two 1s orbitals combine to form 2 molecular orbitals. Orbital 1: Concentrates electron density between the two hydrogen nuclei. Considered constructive interference. Orbital is lower in energy, more stable. Designated a  1s bonding orbital. Orbital 2: Atomic orbitals combine and lead to very little electron density between nuclei. Destructive interference, atomic orbital cancel each other. Higher in energy Designated  1s * antibonding orbital.

35. Order of filling molecular orbitals:  1s  1s *  2s  2s *  2p  2p  2p *  2p * Each  molecular orbital holds 2 electrons; each  orbital holds 4 electrons. Bond order is a measure of the stability of a covalent bond. B.O. = 1/2(number of bonding electrons - number of antibonding electrons) Bond order of 1 = single bond 2 = double bond 3 = triple bond 0 = molecule doesn’t exist

38. Determine the molecular orbital configurations and bond order for the following: C 2 O 2 F 2 Ne 2 Paramagnetism: substances that have one or more unpaired electrons are attracted into a magnetic field and are said to be paramagnetic. Diamagnetism: substances with no unpaired electrons are weakly repelled from a magnetic field and are said to be diamagnetic. Determine whether the above molecules are paramagnetic or diamagnetic.