University:
Bradley University
Course:
CHM 100 | Fundamentals of General Chemistry
Academic year:

2024
Views:

98

Pages:

17
Author:

segolhandworked

Chapter 10 Notes Chemical Bonding II Molecular Shape, Valence Bond Theory, and Molecular Orbital Theory Sections 10.1 – 10.8 I. Artificial Sweeteners: Fooled by Molecular Shape Artificial sweeteners such as aspartame (Nutrasweet) taste sweet like sugar but do not have the calories of sugar. This is because taste and nutritional value are independent properties. The caloric value of food depends on the amount of energy released when the food is metabolized in the body. Many artificial sweeteners are not even metabolized by the body- they just pass straight though. The taste of a food begins with specialized cells on the tongue that detect different molecules of food. These cells are so sensitive that the tongue can detect one molecule of sugar out of thousands of different molecules in a bite of food. We experience certain tastes when molecules of food fit into a special part of a taste receptor on our tongue, much in the same way a key fits into a lock. Artificial sweeteners “trick” our tongues into tasting sweetness because these molecules mimic the molecular shape of the sugar molecule and fit snuggly into our taste receptors. II. VSEPR Theory: The Five Basic Shapes • VSEPR THEORY (VALENCE SHELL ELECTRON PAIR REPULSION): - used to determine the 3- D shape of a molecule - based on the idea that electrons groups (bonds & lone pair e─) repel one another to varying degrees - the combination of repulsions on the central atom of a molecule determines its 3-D shape *The basic idea behind VSEPR is that repulsions between electron groups determine molecular geometry. The preferred geometry is the one in which the electron groups have the maximum separation (and therefore the minimum energy) possible. Preview We will first look at molecular geometries where there are two to six electron groups around a central atom and all of the electron groups are bonding groups (single or double bonds). We will then look at what happens to the shapes when some of the electron groups become lone pair electrons. A. TWO ELECTRONS GROUPS: LINEAR GEOMETRY  Bond Angles: 180° Other examples: SiO2 16 CS2 16 O Si O S C S B. THREE ELECTRONS GROUPS: TRIGONAL PLANAR GEOMETRY Other examples:  Bond Angles: 120° BH3 H SO3 24 6 H O B S H O O In CH2O, the bond angles deviate slightly from 120° because there is more electron density in a double bond than in a single bond. Different electron groups repel each other in slightly different ways. C. FOUR ELECTRONS GROUPS: TETRAHEDRAL GEOMETRY  Bond Angles: 109.5° Other examples: SiCl4 Cl CF4 32 32 Cl F Si C Cl Cl F F F D. FIVE ELECTRONS GROUPS: TRIGONAL BIPYRAMIDAL GEOMETRY  Bond Angles: 90°, 120° Other examples: SOF4 ClO2F3 40 40 O O F F S Cl F F F F F O E. SIX ELECTRONS GROUPS: OCTAHEDRAL GEOMETRY  Bond Angles: 90° Other examples: SCl6 48 Cl Cl Cl S Cl Cl Cl III. VSEPR Theory: The Effect of Lone Pairs Preview We will now look at molecular geometries where there are four to six electron groups around a central atom and some of the electron groups are lone pairs. Keep in mind that these shapes are all variations of the geometries from the last section, but now one or more bonding groups have been replaced with lone pairs. The difference in shapes results from the fact that lone pair electrons repel other lone pair and bonding electrons to a greater extent than bonding electrons repel one another. THE ORDER OF ELECTRON PAIR REPULSION: lone pair - lone pair > lone pair- bonding pair > bonding pair- bonding pair Most repulsive Least repulsive * Want to be as far away from each other as possible THE DIFFERENCE BETWEEN:  Electron Geometry: The geometrical arrangement of electron groups Molecular Geometry: The geometrical arrangement of the atoms • In this section, don’t get confused between electron and molecular geometries. • In the last section, the electron and molecular geometries were the same (same name for each). • In this section, they are different. • There are only 5 electron geometries, but there are 11 molecular geometries. A. FOUR ELECTRON GROUPS WITH LONE PAIRS (TETRAHEDRAL ELECTRON GEOMETRY) TRIGONAL PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR  Bond Angles: 107° Other examples: PF3 26 NCl3 26 P F N F F Cl Cl Cl A. FOUR ELECTRON GROUPS WITH LONE PAIRS (TETRAHEDRAL ELECTRON GEOMETRY) BENT (MOLEC. GEO.): TWO LONE PAIRS  Bond Angles: 104.5° Other examples: SCl2 18 Cl H2S S Cl H 8 S H Summarizing Tetrahedral Electron Geometries B. FIVE ELECTRON GROUPS WITH LONE PAIRS (BIPYRAMIDAL ELECTRON GEOMETRY) SEESAW (MOLEC. GEO.): ONE LONE PAIR  Bond Angles: 90°, 120° Other examples: SeCl4 IOF3 34 34 Cl O Cl F I Se Cl Lone pair must go equatorial Cl F F B. FIVE ELECTRON GROUPS WITH LONE PAIRS (BIPYRAMIDAL ELECTRON GEOMETRY) T-SHAPED OR LINEAR (MOLEC. GEO.): TWO OR THREE LONE PAIRS  Bond Angles: 90°, 120° Other examples: I3 - ClF3 28 F Cl I F F Lone pairs must go equatorial I I 28 C. SIX ELECTRON GROUPS WITH LONE PAIRS (OCTAHEDRAL ELECTRON GEOMETRY) SQUARE PYRAMIDAL (MOLEC. GEO.): ONE LONE PAIR SQUARE PLANAR (MOLEC. GEO.): TWO LONE PAIRS  Bond Angles: 90° Other examples: BrF4- XeOF4 42 36 O F F F Xe F F Br F Lone pairs must go axial F F

Molecular Shape, Valence Bond Theory, and Molecular Orbital Theory

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