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Chapter 10 Molecular Shapes and Valence Bond Theory 10.1 Artificial Sweeteners: Fooled by Molecular Shape (Suggested Reading) 10.2 VSEPR Theory: The Five Basic Shapes 10.3 VSEPR Theory: The Effect of Lone Pairs 10.4 VSEPR Theory: Predicting Molecular Geometries 10.5 Molecular Shape and Polarity 10.6 Valence Bond Theory: Orbital Overlap as a Chemical Bond 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals (sp, sp2 and sp3 only) 1 Lewis dot structure Molecular Shape • Lewis dot structures only gives us an idea of the electron distribution in the species. There is NO IDEA about the molecular geometry, which depends on the relative position of terminal atoms around the central atom. • We will connect the electron distribution in a Lewis dot structure → molecular geometry by using the ValenceShell Electron-Pair Repulsion (VSEPR) theory. VSEPR = Groups of electrons repel each other, ending up as far from each other as physically possible. 2 1 VSEPR Model • Although the Theory states “repulsion of ELECTRON PAIRS” ………..it is actually repulsion of ELECTRON GROUPS because lone pair, single, double and triple bond pairs are treated as ONE PAIR of electrons in VSEPR theory That is: A lone pair is ONE GROUP of electrons A single bond is ONE GROUP of electrons A double bond is ONE GROUP of electrons A triple bond is ONE GROUP of electrons 3 Electron distribution vs. geometry Electron distribution Molecular geometry Looks at “shape” of electron group distribution Looks at “shape” of nuclear positions around the central atom INCLUDES lone pairs No terminal nuclei on lone pairs means we IGNORE (“can’t see”) all lone pairs. Bonding pair of eLone pair of e- 4 2 Electron distribution vs. Geometry If the central atom has no lone pairs on it, then the electron group distribution and the molecular geometry are the same ! 5 Molecular Geometries Examples of geometries of molecules with no lone pairs around central atom and bond angles Lewis Structure For CH4 6 3 Molecular Geometries Its molecular shape Lewis structure for PCl5 Lewis structure for SF6 7 The Effect of Lone Pairs on Shape • Lone pair groups “occupy more space” on the central atom. – because their electron density is exclusively on the central atom rather than shared like bonding electron groups • relative sizes of repulsive force interactions are: • Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • This affects the bond angles, making them smaller than expected. 8 4 The Effect of Lone Pairs on Shape IF NO LONE PAIRS -The molecule’s shape will be one of the basic molecular geometries if all the electron groups are bonds (see slide 5) LONE PAIRS AROUND CENTRAL -Molecules with lone pairs will have distorted bond angles but the shape will be a derivative of one of the basic shapes (See Table 10.1(lone pair column) for examples) 9 Electron and Molecular Geometries 10 5 Electron and Molecular Geometries Electron groups 11 Electron and Molecular Geometries Electron groups 12 6 The Effect of Lone Pairs on Shape Detailed examples of 4 molecules (1) When there are three electron groups around the central and one of them is a lone pair, the resulting molecular shape is called a bent shape. The bond angle is <120°. SO2 Example of 2 bonding and one lone pair <1200 Bent shape 13 The Effect of Lone Pairs on Shape (2) When there are four electron groups around the central and one of them is a lone pair Lewis structure of NH3 <1090 Example of 3 bonding and 1 lone pair 14 7 The Effect of Lone Pairs on Shape (3) When there are five electron groups around the central and one of them is a lone pair A “seesaw” shape Example of 4 bonding And one lone pair 15 The Effect of Lone Pairs on Shape (4) When there are six electron groups around the central and one of them is a lone pair All <900 Lewis structure of BrF5 Example of 5 bonding And one lone pair 16 8 Bond Angles You also need to know the bond angles of all the shapes for: 1) The basic molecular geometries AND 2) the geometries of molecules having lone pairs (Table 10.1) 17 Figuring out Molecular Shapes 1. Draw the Lewis dot structure 2. Determine the number of electron groups on the central atom to get electron geometry (see Slide 5). If no lone pairs around central atom then you have the molecule’s molecular geometry. 3. If central atom has lone pairs, use the number of bonded and lone pairs and the arrangement (Table 10.1) to determine resulting molecular geometry. 4. Draw the 3 D structure as best you can. 3. Determine bond angles. 18 9 Drawing 3D shapes • By convention, the central atom is put in the plane of the paper. • Put as many other atoms as possible in the same plane and indicate with a straight line. • For atoms in front of the plane, use a solid wedge. • For atoms behind the plane, use a hashed wedge 19 Practice: Draw the 1) Lewis structures 2) the molecular shapes ( 3D structure) around their central atom and 3) the bond angles for the following molecules? (a) (b) (c) (d) (e) H2O PF5 SeCl4 KrCl2 IF5 20 10 Dipole Moment and Molecular Shape • If there are polar covalent bonds in a molecule, the molecule MAY OR MAY NOT have a permanent dipole moment. • A permanent dipole moment means that there is a partially negative and a partially positive site that is permanent. • To determine if a molecule has a permanent dipole moment, we add together the vectors of all the polar covalent bonds (and their dipole moments.) 21 Molecular Dipole Moment • A simple permanent dipole is HCl •• H - Cl : •• δ+ δ− Dipole moment = 3.34 D •• It has a polar covalent bond. Since H - Cl : there is •• only one bond, this one vector of charge describes the permanent dipole. 22 11 Molecular Dipole Moment • Water has two polar covalent bonds and two dipoles • The permanent dipole moment in water can be seen by adding together the charge separation vectors of the two polar covalent O-H bonds. Dipole moment = 1.94 D 23 Molecular Dipole Moment But a molecule with more than one polar bond MIGHT NOT have a permanent dipole moment . If there is symmetry of the (identical) polar bonds the resultant vector sums may add up to zero. An example is carbon dioxide CO2 24 12 Predicting dipole moments from Geometry Note: 25 Predicting dipole moments from Geometry Practice—Decide whether the following molecules are polar, given the EN values. EN values O = 3.5 N = 3.0 Cl = 3.0 S = 2.5 26 13 Molecular Properties from molecular dipole moments • As we will see in next Chapter (section 11.3) molecules with partial positive and negative charges will attract the opposite regions on other molecules of the same type. Such intermolecular forces affect the molecular properties of the compound, i.e., It is a liquid or a solid. It will also affect the compound’s boiling point. Example: two isomers of C2H2Cl2 cis-1,2-dichloroethane trans-1,2-dichloroethane boiling point 60oC 48oC 27 Explain why the b.p. are different by drawing the dipole moments of each. 10.6 Valence Bond Theory (Orbital overlap) Covalent bonds form between atoms when: 1. Orbitals in the atoms overlap to create molecular bonding orbitals. 2. Each molecular bonding orbital has no more than 2 electrons in it. ALSO…. bond formation occurs between two atomic orbitals containing one electron and …..Covalent bonds are strongest when there is maximum orbital overlap between atomic orbitals. 28 14 Valence Bond Theory (Orbital overlap) Here are some favourable atomic orbital overlaps for H2 and HCl H 1s ↑ H 1s ↑ 29 10.7 Hybridization of atomic orbitals • But there are problems that arise from this simple theory. The number of partially filled or empty atomic orbitals did not always predict the number of bonds or orientation of bonds. For Carbon whose valence atomic orbital is = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart in CH4. • To adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridize (mix) before bonding took place. 30 15 Unhybridized C orbitals predict the wrong bonding and geometry Hybridization of C is to mix all the 2s and 2p orbitals to get four equal orbitals that point to the corners of a tetrahedron. 31 Hybridization • Many atoms hybridize their orbitals to maximize bonding. – Hybridizing is mixing different types of orbitals to make a new set of degenerate orbitals. – sp, sp2, sp3, sp3d, sp3d2 – more bonds = more full orbitals = more stability • Same types of atom can have different hybridizations depending on the compound. C = sp, sp2, sp3 32 16 Carbon Hybridizations Unhybridized ↑↓ ↑ ↑ 2p 2s sp hybridized ↑ ↑ ↑ 2sp ↑ 2p sp2 hybridized ↑ ↑ ↑ 2sp2 sp3 hybridized ↑ ↑ ↑ ↑ 2p ↑ 2sp3 33 Hybridization • The number of standard atomic orbitals combined equals the number of hybrid orbitals formed. – H cannot hybridize! . • The number and type of standard atomic orbitals combined determine the shape of the hybrid orbitals. • The particular kind of hybridization that occurs is the one that yields the lowest overall energy for the molecule. 34 17 Hybridization Orbital Diagram of the sp3 Hybridization of C 35 Formation of sp3 Hybrid Orbitals 36 18 Hybridization sp3 Hybridized Atoms: Orbital Diagrams • Place electrons into hybrid and unhybridized valence orbitals as if all the orbitals have equal energy. • Lone pairs generally occupy hybrid orbitals. Unhybridized atom 2s ↑↓ 2s ↑ ↑ 2p C ↑ ↑ ↑ ↑ 2p N ↑ ↑ ↑ ↑ 2sp3 ↑ ↑↓ sp3 hybridized atom ↑ ↑ ↑ 2sp3 37 Hybridization Methane Formation with sp3 C Ammonia Formation with sp3 N 38 19 Other Types of Hybrid Orbitals • A total of n atomic orbitals combine to give n hybrid orbitals of a given kind 39 Resulting Shapes of Hybrid Orbitals 40 20 Determining hybrid orbital diagrams 1. 2. 3. 4. Draw the Lewis dot structure Use VSEPR theory to predict electron group arrangement Use Table on slide 38 to determine what hybrid orbitals have the same arrangement Create the hybrid orbital diagram based on changing the ground state diagram of the central atom Practice: Describe the bonding of BF3 in terms of hybrid orbital theory. 41 Types of Bonds • A sigma (σ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei. -between s-to-s, hybrid-to-hybrid, s-to-hybrid, axis p • A pi (π) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei. – between unhybridized parallel p orbitals 42 21 Types of Bonds 43 Orbital Diagrams with Hybridization • “Overlap” between a hybrid orbital on one atom and a hybrid or nonhybridized orbital on another atom results in a σ bond. • “Overlap” between unhybridized p orbitals on bonded atoms results in a π bond. 44 22 An example of a Double Bond from hybridization H2C=CH2 Sp2 orbitals Double bond using 1 sp2 and the 2p 2p orbitals In ethene (C2H4) the hybrid orbital is an sp2 (not sp3) 45 An example of a Triple Bond C2H2 (ethyne) HC≡CH using sp hybridized Using 2p atomic orbitals 46 23 Example of sp2 Hybridized C and O in CH2O Formaldehyde pC ↑ π σ sp2 C ↑ ↑ ↑ ↑ ↑↓ ↑↓ sp2 O σ ↑ σ pO ↑ ↑ 1s H 1s H 47 Practice—Draw the orbital diagram for hybridization of B and O atom. How many σ the and π bonds would you expect each to form? sp2 Unhybridized atom ↑↓ 2s B ↑ 2p O ↑↓ 2s ↑↓ ↑ ↑ 2p 48 24 HCN Orbital Diagram-sp hybridized ↑ pC sp C ↑ 2π ↑ ↑ σ ↑ ↑ pN ↑ ↑↓ sp N s ↑ 1s H 49 sp3d Hybridized -Orbital Diagrams Unhybridized atom ↑↓ 3s ↑↓ 3s ↑ ↑ ↑ 3p ↑↓ ↑ ↑ 3p sp3d hybridized atom P 3d ↑ ↑ ↑ ↑ 3sp3d S ↑↓ ↑ 3d ↑ ↑ ↑ ↑ 3sp3d (nonhybridizing d orbitals not shown) 50 25 SOF4 Orbital Diagram ↑ dS ↑ ↑ σ σ σ ↑ ↑ σ ↑ ↑ sp3d S ↑ ↑ π σ pO ↑ ↑ ↑↓ ↑↓ sp2 O ↑ 2p F 2p F 2p F 2p F 51 26