Unit 28 Molecular Geometry

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1 Unit 28 Molecular Geometry There are two concepts in the study of molecular geometry. One is called the Valence Shell Electron Pair Repulsion (VSEPR) model. The other is electron orbital hybridization. Both are based on the ability of the student to draw Lewis dot structures of atoms and molecules. We will start with a quick review. Lewis Structures. Lewis structures are based on the concept that all atoms want filled orbitals. Usually atoms want to duplicate the 8 valence (outer) electrons of the noble gases. This is commonly called the Octet Rule. This can be done in two ways. One is to completely exchange electrons, which is done in the formation of ionic compounds and in oxidation and reduction reactions. Another is to share electrons in covalent bonds. In this experiment we will look at shared e - in covalent bonds. To determine the Lewis structure for a molecule you can follow the steps below for HNO3: 1. Determine the basic structure of the atom by selecting the central atom (usually the atom with the highest oxidation number that occurs alone) and placing the other atoms around it. For example: O N O O H 2. Determine the number of valence electrons for each atom. This is done from the group (column) the atom occupies in the periodic table. Group 1 has 1 valence electron, group 2 has 2 valence electrons, etc. Sketch the valence electrons around each atom in pencil. For N, valence electrons = 5 For O, valence electrons = 6 For H, valence electrons = 1 H.. : O :.... : N : O :.. : O :..

2 3. By moving the valence electrons around the source atom, arrange the atoms so that there are 8 original electrons and/or shared electrons around each atom. There can be no more than 8 electrons around each atom except hydrogen, which wants only two electrons. Sharing of 4 electrons constitutes a double bond. Sharing of 6 electrons constitute a triple bond... : O :.... H : N : O :.. : O :.. 4. The lone pairs around the central atom contribute to the geometry of the molecule. VSEPR model. The VSEPR model is based on the repulsion of electron pairs of molecules that have a central atom. The central atom defines the shape of the molecule. The VSEPR description begins with a central atom "A" followed by the number of atoms attached to the central atom, each designated as an "X". The final designation is the number of lone pairs of electrons each described as "E". For example, water (H2O) has oxygen as a central atom (A) with two hydrogens attached (X2) and two lone pair (E2) to give a VSEPR designation of AX2E2. Ammonia (NH3) has a nitrogen central atom, three hydrogen appendages, and one lone pair to give AX3E. Hybridization. When atomic orbitals combine in covalent bonds, they change their nature and shape. The shared electrons now form a molecular orbital that is a hybrid of the individual atomic orbitals. These hybrids are affected by space and electronic charge repulsions of the electrons. These new orbitals are all equal in shape and size but are made up of atomic orbitals that are not equal in shape or size. The most common is the hybrid orbitals that contain the stable 8 valence electrons similar to the noble gases. These 8 electrons form 4 electron-pair orbitals made from an "s" atomic orbital and 3 "p" atomic orbitals. They are called sp 3 hybrid molecular orbitals. They are all identical even though they are made from "s" and "p" electrons. 266

3 When describing the hybridization of the molecular orbitals of a molecule, like in the VESPR model, you start with a central atom. Unlike the VSEPR model, it does not matter whether other atoms or just electron pairs are attached to the central atom. You want to look at all the electron pairs whether or not they are bonded to another atom. The maximum number of bonding pairs is 6. The minimum is 1. To determine the hybridization, count the number of electron pairs extending from the central atom. Then follow this scheme: atomic orbitals s p p p d d d d d # of e- pairs Then you combine like atomic orbitals to give the following possibilities: S sp sp 2 sp 3 sp 3 d sp 3 d 2 You will notice that you cannot have 7, 8, or 9 electron pairs in hybridization even though it follows the pattern above. In these cases there would be too many electrons to spatially fit around a central atom. On the next page is a chart showing the common molecular shapes, the VSEPR designation and the hybridization. 267

4 ELECTRON PAIRS Total Bonding Lone Shape VSEPR Hybridization Example linear AX2 sp BeF trigonal AX3 sp 2 BF3 planar bent AX2E sp 2 SO tetrahedral AX4 sp 3 CH trigonal AX3E sp 3 NH3 pyramidal bent AX2E2 sp 3 H2O ELECTRON PAIRS Total Bonding Lone Shape VSEPR Hybridization Example trigonal AX5 sp 3 d PCl5 bipyramidal seesaw AX4E sp 3 d SF T-shaped AX3E2 sp 3 d ClF linear AX2E3 sp 3 d XeF octahedral AX6 sp 3 d 2 SF square AX5E sp 3 d 2 IF5 pyramidal square AX4E2 sp 3 d 2 XeF4 planar 268

5 Organic Molecular Geometry. The element carbon, with its 4 valence electrons and small size provide a wide variety of molecular combinations with varying shapes. In order to achieve the octet rule, double and triple bonds can be formed. We will use molecular models to explore the shapes of some organic compounds and the effect of applying double and triple bonds. When two different compounds have the same molecular formula but different structure, we call them geometric or structural isomers. A good example is butane and isobutane. The term "iso" means branched. Some isomers are so close that they differ only in the fact that one is the mirror image of the other. Mirror images are similar but cannot be superimposed onto each other. These mirror image isomers are called enantiomers, optical isomers, or stereoisomers. The difference between two is that one enantiomer will rotate plane polarized light in one direction and the other will rotate plane polarized light in the opposite direction. The enantiomer that rotates the light to the right is called the "dextrorotatory" isomer and the one that rotates light to the left is the "levorotatory" isomer. This is abbreviated in the compound's name as d-glucose and l-glucose for the glucose stereoisomers. A chiral carbon, or a carbon with 4 completely different appendages or branches will determine if a mirror image exists. When looking for 4 completely different appendages you must look beyond the closest atom to the whole branch. For every chiral carbon that you find, a mirror image or another enantiomer exists. This topic will be explored in greater depth in Unit 29. Some isomers occur because 2 branches either occur on the same side or in close proximity to each other (cis isomer) or they occur on opposite sides of the molecule from each other (trans isomer). 269

6 270

7 Prelab Exercises for Unit 28 Name Section Date 1. What is the Lewis Structure, shape, VSEPR designation and hybridization associated with the following compounds? Lewis Structure Shape VSEPR Hybridization CCl4 H2S NH3 SO4-2 CO2 NO3 - (over) 271

8 2. Draw 3 examples of compounds with chiral carbons along with their enantiomers. 272

9 Lab Report for Unit 28 Name Section Date Using the molecular models, make the following molecules, draw their Lewis Structure, sketch their shape, and give any other requested information. 1. carbon dioxide Sketch & Lewis Structure: shape 2. hydrogen cyanide (HCN) Sketch & Lewis Structure: shape 3. nitrous acid (HNO2) Sketch & Lewis Structure: shape 273

10 4. silicon tetrachloride (SiCl4) Sketch & Lewis Structure: shape 5. tellurium tetrachloride (TeCl4) Sketch & Lewis Structure: shape 6. Iodine trichloride (ICl3) Sketch & Lewis Structure: shape 7. sulfur hexafluoride (SF6) Sketch & Lewis Structure: shape 274

11 8. Propane (C3H8) Sketch & Lewis Structure: 9. Butane and geometric isomers (C4H10) Sketches & Lewis Structures: chlorobutane and any optical isomers (C4H9Cl) Sketches & Lewis Structures: 275

12 11. 2,3-dichlorobutane and any optical isomers (C4H8Cl2) Sketches & Lewis Structures: 12. ethylene (C2H4) Sketch and Lewis Structure: 13. cis-dibromoethylene (C2H2Br2) and trans-dibromoethylene Sketches & Lewis Structure: 276

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