Transcript Chapter 9
Chapter 9 Molecular Geometry Introduction 1. Lewis Structures help us understand the compositions of molecules & their covalent bonds, but not their overall shapes. 2. The properties of a substance largely depend on the shape & size of its molecules, together with the strength & polarity of its bonds. 3. Examples: Taxol, smell, vision VSEPR Theory Valence-shell electron-pair repulsion 1. The overall shape of a molecule is determined by its bond angles. 2. The VSEPR Model is used to predict Molecular Shapes or Molecular Geometries. 3. The VSEPR Theory assumes that each atom in a molecule will be positioned so that there is minimal repulsion between the valence electrons of that atom. Five Basic Shapes 1. Linear 2. Trigonal Planar 3. Tetrahedral 4. Trigonal Bipyramidal 5. Octahedral Chart to Memorize for Basic Shapes Bond Angle(s) #Electron-Pairs Hybridization 180 2 sp 120 3 sp2 109.5 4 sp 3 90 & 120 5 sp3d 90 6 sp 3d2 Using VSEPR Model to Predict Shapes 1. Draw the Lewis dot structure to determine the total # of electron pairs around the central atom. 2. Multiple bonds (double and triple) count as one. 3. The total number of bonding and nonbonding electron pairs determines the geometry of the electron pairs (one of the 5 basic shapes). 4. Then use bonding electron pairs only to determine the molecular geometry (actual shape of molecule). Special Rules 1. Lone pairs effect geometry more than bonding pairs. 2. NH3 has one lone pair: reduces angle from 109.5 to 107 3. H2O with two lone pairs: reduces angle from 109.5 to 105. 4. Multiple bonds affect geometry more than single bonds 5. H2C=O (116 instead of 120) 6. H2C=CH2 (117 instead of 120) 7. Lone pairs occupy the axial (middle) position for trigonal bipyramidal structures. Practice Determine the geometry of the following: BeCl2 CO2 BF3 O3 SO2 CH4 PCl3 H2O PCl5 SF4 ClF3 XeF2 SF6 IF5 XeF4 Consult the next 3 pages to help you. Linear BeCl2 valence e- = Be 2 + (2 x 7) = 16e- .. Cl .. .. .. .. Cl .. 180o Two Electron Pairs = Linear Molecule 180o linear CO2 valence e- = two C .. O .. valence pairs on C .. .. .. O .. 4 + (2 x 6) = 16e- .. O .. C .. O .. ignore double bonds single and double bonds same molecular geometry linear molecular shape linear 120o trigonal planar SO2 .. .. O : S .. O .. .. .. O .. = 18e- : .. O .. (2 x 6) .. three S 6+ .. .. .. O .. : valence e- = S valence pairs on S two bonding pairs one lone pair molecular geometry molecular shape < 120o trigonal bent .. .. O tetrahedral 109.5o CH4 valence e- = 4+ (4 x 1) = 8eH four valence pairs on C H 109.5o C H H molecular geometry molecular shape tetrahedral tetrahedral tetrahedral 109.5o NH3 valence e- = 5+ (3 x 1) = 8e- : four H valence pairs on N N H three bonding pairs H one lone pair molecular geometry molecular shape < 109.5o tetrahedral trigonal pyramid bipyramidal 120o and 1800 PCl5 5+ (5 x 7) Cl = 40e- five valence pairs on P P 90o molecular geometry 120o molecular shape bipyramidal bipyramidal .. .. 180o .. .. .. Cl .. Cl .. .. .. Cl .. .. Cl .. .. .. valence e- = .. bipyramidal 120o and 1800 SF4 valence e- = 6+ (4 x 7) = 34e- : five .. four bonding pairs .. one lone pair S < 180o molecular geometry molecular shape .. F .. .. F .. .. .. .. F .. .. F .. valence pairs on S bipyramidal seesaw bipyramidal 120o and 1800 ClF3 valence e- = 7+ (3 x 7) = 28e- : five .. three bonding pairs two lone pair molecular geometry molecular shape 90o 180o Cl .. F .. .. F .. .. .. .. F .. valence pairs on Cl bipyramidal T 120o and 1800 bipyramidal ICl2valence e- = 7+ (2 x 7) + e- = 22e- : five .. I two bonding pairs .. Cl .. three lone pair on I molecular geometry molecular shape bipyramidal linear .. .. Cl .. valence pairs on I octahedral 90o BrF5 valence e- = 7+ (5 x 7) = 42e- : .. six valence pairs on Br Br .. .. five bonding pairs molecular geometry molecular shape .. .. .. F one lone pair .. .. .. F .. F .. .. .. .. F .. F octahedral square pyramidal octahedral 90o XeF4 valence e- = 8+ (4 x 7) = 36e- : valence pairs on Xe two lone pair molecular geometry molecular shape .. .. four bonding pairs Xe : six .. F .. .. F .. .. .. .. F .. .. F .. octahedral square planar : .. .. : O F B S .. .. : H : C S H O F O H O : : .. S :F : .. O .. S C : : .. .. O .. Be .. .. H .. .. .. O .. Cl .. .. O .. .. .. .. O .. Cl .. H C .. .. O : N H .. O .. .. .. : .. O .. H : .. .. .. .. .. .. .. Cl .. I .. Cl .. .. .. .. .. F .. F .. Cl .. .. .. .. .. .. .. F .. .. : : : .. S .. F .. .. F .. .. S F .. F .. .. F .. .. F .. .. F .. .. .. .. : .. F .. .. F .. .. P .. Cl .. .. Cl .. .. F .. H .. O H .. Cl .. .. Cl .. .. .. : Cl .. F .. .. : Br .. .. .. F .. .. F .. .. .. .. .. .. Xe .. F .. .. F .. .. F .. .. F .. : : .. .. F .. .. F .. Valence Bond Theory 1. Valence Bond Theory explains why molecules have the shapes they do based on a concept called hybridization. 2. Hybridization - A mixture of two or more atomic orbitals. http://www.youtube.com/watch?v=PrNbhuB9W44&feature=related 10 Minutes Valence Bond Theory & Hybridization 1. In 1931, Linus Pauling, proposed that the outermost orbitals of an atom could be combined to form hybrid atomic orbitals. – Sigma bond (s) - The end-to-end overlapping of two orbitals. – Pi bond (p) - The side-to-side overlapping of two p orbitals. • Single bonds are made up of one sigma bond. • Double bonds are made up of one sigma bond and one pi bond. • Triple bonds are made up of one sigma bond and two pi bonds. Sigma Bonds The electron density is concentrated on the internuclear line. More Sigma Bonds Pi Bonds The electron density is concentrated above & below the internuclear line. Molecular Orbital Theory MO Theory is used to predict: 1) whether or not a molecule exists 2) its bond type (single, double, or triple) 3) its bond strength (triple are strongest) and 4) some of its properties (paramagnetic or diamagnetic). MO Diagrams & Magnetic Properties Paramagnetism – Molecules with one or more UNPAIRED electrons in their MO diagram are attracted into a magnetic field. Diamagnetism – Molecules with PAIRED electrons in their MO diagram are weakly repelled from a magnetic field Molecular Orbitals 1. Molecular orbitals, like atomic orbitals, have a definite shape and hold up to 2 electrons of opposite spin. 2. When 2 atomic orbitals combine & overlap, they form 2 molecular orbitals. 3. One of these molecular orbitals in a BONDING orbital and the other is an ANTIBONDING orbital. Bonding molecular orbital - Electrons in this orbital spend most of their time in the region directly between the two nuclei. Antibonding molecular orbital - Electrons placed in this orbital spend most of their time away from the region between the two nuclei. Shown below are a sigma (s) bonding molecular orbital and a sigma antibonding (s *), molecular orbital. * = Antibond (a) Shows sigma MOs (b) & (c) Shows pi MOs Molecular Orbital Diagram Small 2s-2p Interaction Use for Oxygen, Fluorine, & Neon (Reverse order for o 2p and p 2p for all other others) Bond Order Bond order Type of Bond 0 Molecule doesn’t exist 1 Single bond 2 Double bond 3 Triple bond Molecular Orbital Diagram Use with the Bond Order Formula to determine if a molecule exists, its type of bond, & whether it is paramagnetic or diamagnetic. (Small 2s-2p Interaction) Use for Oxygen, Fluorine, & Neon (Large 2s-2p Interaction) Use for Boron, Carbon, Nitrogen Molecular Orbital Diagrams - Steps to determine bond type 1) Sum valence electrons for both atoms 2) Place arrows (represent electrons) in center of the diagram 3) Determine Bond Order = (# bonding e- - # antibonding e-)/2 4) ID Bond Type (1 = single, 2 = double, 3 = triple) 5) ID as paramagnetic (unpaired e- = attract magnet) or diamagnetic (all paired e- = repel magnet) Use for Oxygen, Fluorine, & Neon Use for Boron, Carbon, Nitrogen • http://www1.teachertube.com/viewVideo.ph p?video_id=57124 • A Little MO Theory • 14 minutes