Chapter 10 Molecular Shapes and Valence Bond Theory

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1 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, sp 2 and sp 3 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 Valence- Shell Electron-Pair Repulsion (VSEPR) theory. VSEPR = Groups of electrons repel each other, ending up as far from each other as physically possible. 2 1

2 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 Looks at shape of electron group distribution INCLUDES lone pairs Molecular geometry Looks at shape of nuclear positions around the central atom No terminal nuclei on lone pairs means we IGNORE ( can t see ) all lone pairs. Bonding pair of e - Lone pair of e - 4 2

3 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 CH 4 6 3

4 Molecular Geometries Its molecular shape Lewis structure for PCl 5 Lewis structure for SF 6 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

5 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

6 Electron and Molecular Geometries Electron groups 11 Electron and Molecular Geometries Electron groups 12 6

7 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. SO 2 Example of 2 bonding and one lone pair <120 0 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 NH 3 <109 0 Example of 3 bonding and 1 lone pair 14 7

8 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 <90 0 Lewis structure of BrF 5 Example of 5 bonding And one lone pair 16 8

9 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

10 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) H 2 O (b) PF 5 (c) SeCl 4 (d) KrCl 2 (e) IF

11 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. H - Cl Since there is only one bond, this one vector of charge describes the permanent dipole. : 22 11

12 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 CO

13 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 =

14 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 C 2 H 2 Cl 2 boiling point cis-1,2-dichloroethane 60 o C trans-1,2-dichloroethane 48 o C 27 Explain why the b.p. are different by drawing the dipole moments of each 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

15 Valence Bond Theory (Orbital overlap) Here are some favourable atomic orbital overlaps for H 2 and HCl H 1s H 1s 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 = 2s 2 2p x1 2p y1 2p z0 would predict two or three bonds that are 90 apart, rather than four bonds that are apart in CH 4. To adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridize (mix) before bonding took place

16 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, sp 2, sp 3, sp 3 d, sp 3 d 2 more bonds = more full orbitals = more stability Same types of atom can have different hybridizations depending on the compound. C = sp, sp 2, sp

17 Carbon Hybridizations Unhybridized 2s 2p sp hybridized 2sp 2p sp 2 hybridized sp 3 hybridized 2sp 2 2p 2sp 3 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

18 Hybridization Orbital Diagram of the sp 3 Hybridization of C 35 Formation of sp 3 Hybrid Orbitals 36 18

19 Hybridization sp 3 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 2p C sp 3 hybridized atom 2sp 3 2s 2p N 2sp 3 37 Hybridization Methane Formation with sp 3 C Ammonia Formation with sp 3 N 38 19

20 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

21 Determining hybrid orbital diagrams 1. Draw the Lewis dot structure 2. Use VSEPR theory to predict electron group arrangement 3. Use Table on slide 38 to determine what hybrid orbitals have the same arrangement 4. Create the hybrid orbital diagram based on changing the ground state diagram of the central atom Practice: Describe the bonding of BF 3 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

22 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

23 An example of a Double Bond from hybridization H 2 C=CH 2 Sp 2 orbitals Double bond using 1 sp 2 and the 2p 2p orbitals In ethene (C 2 H 4 ) the hybrid orbital is an sp 2 (not sp 3 ) 45 An example of a Triple Bond C 2 H 2 (ethyne) HC CH using sp hybridized Using 2p atomic orbitals 46 23

24 Example of sp 2 Hybridized C and O in CH 2 O Formaldehyde p C sp 2 C π σ p O sp 2 O σ σ 1s H 1s H 47 Practice Draw the orbital diagram for the sp 2 hybridization of B and O atom. How many σ and π bonds would you expect each to form? Unhybridized atom 2s 2p 2s 2p B O 48 24

25 HCN Orbital Diagram-sp hybridized sp C p C 2 π σ p N sp N s 1s H 49 sp 3 d Hybridized -Orbital Diagrams Unhybridized atom sp 3 d hybridized atom 3s 3p 3d P 3sp 3 d 3s 3p 3d S 3sp 3 d (nonhybridizing d orbitals not shown) 50 25

26 SOF 4 Orbital Diagram d S sp 3 d S π σ p O sp 2 O σ σ σ σ 2p F 2p F 2p F 2p F 51 26

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