# Shapes of Molecules. AX m E n. A. Molecular Geometry

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1 Shapes of Molecules A. Molecular Geometry Lewis structures provide us with the number and types of bonds around a central atom, as well as any NB electron pairs. They do not tell us the 3-D structure of the molecule. C C 4 as drawn conveys no 3-D information (bonds appear like they are 90 apart) The Valence Shell Electron Pair Repulsion Theory (VSEPR), developed in part by Ron Gillespie at McMaster in 1957, allows us to predict 3-D shape. This important Canadian innovation is found worldwide in any intro chem course. VSEPR theory has four assumptions 1. Electrons, in pairs, are placed in the valence shell of the central atom 2. Both bonding and non-bonding (NB) pairs are included 3. Electron pairs repel each other maximum separation. 4. NB pairs repel more strongly than bonding pairs, because the NB pairs are attracted to only one nucleus To be able to use VSEPR theory to predict shapes, the molecule first needs to be drawn in its Lewis structure. 1 VSEPR theory uses the AXE notation (m and n are integers), where m + n = number of regions of electron density (sometimes also called number of charge clouds). AX m E n 1. Molecules with no NB pairs and only single bonds We will first consider molecules that do not have multiple bonds nor NB pairs around the central atom (n = 0). Example: BeCl 2 o Molecule is linear (180 ) Example: B 3 2 o Molecule is trigonal (or triangular) planar (120 )

2 Example: C 4 o Molecule is tetrahedral (109.5 ) Example: P Molecules with 1 NB pairs and only single bonds The geometry of the regions of electron density is roughly the same as what we see when no NB pairs are involved. owever, the shape of the molecule is determined by looking at only the bonding pairs, NOT the non-bonding pairs. Example: N 3 4 o Molecule is trigonal bipyramidal (90 and 120 ). There are three X atoms in a planar triangle and two axial atoms, one above and one below the central atom. Example: S 6 o There are four regions of electron density (m + n = 4), and the electronic arrangement is still tetrahedral. o Yet, the shape of the molecule (look at bonding pairs only) is trigonal pyramidal. Angles < (why?) Example: 2 O o Molecule is octahedral (all 90 ) o Electronic arrangement is tetrahedral. Shape of molecule = bent

3 Example: Xe 2 (or I 3 - ) 5 Example: Xe o Electronic arrangement is trigonal bipyramid. Shape of molecule = linear Example: S 4 o Electronic arrangement is trigonal bipyramidal. Shape of molecule = T-shaped Why is this better than triganol planar? Xe o Electronic arrangement is trigonal bipyramidal. Shape of molecule = see saw Why is this better than triganol pyramid? S Example: Xe 5 + o Electronic arrangement is octahedral. Shape of molecule = Square pyramidal

4 7 8 Example: ICl 4 3. Molecules with double or triple bonds These are quite simple: treat multiple bonds as single bonds, and the AXE system still works. i.e. a multiple bond is still considered to be one region of electron density. Examples: predict the shapes of CO 2, C 2 4, and NO 2 o Electronic arrangement is octahedral. Shape of molecule = Square planar Why is this better than? Cl I Cl Cl Cl Example: the amino acid alanine 3 C o There can be more than one central atom C N 2 O C O

5 9 10 Summary

6 B. Molecular Polarity Recall that electronegativity is a relative measurement of an atom s ability to attract a bonding electron pair to itself. 11 Note: The overall polarity of a molecule depends on its geometry and the presence of polar bonds. o The presence of polar bonds does not necessarily imply that the molecule is polar. This is because the bond dipole moments can cancel out (no net dipole moment). o Think of the bond dipoles as vectors, and if the sum of all of the vectors is zero, the molecule is not polar. 12 Examples: Differences in electronegativity between two covalently bonded atoms result in a polar covalent bond, and atoms have a partial negative or partial positive charge. CO 2 2 O These bonds are said to have dipole moments, which can be experimentally quantified. If placed in an electric field, the molecule will rotate and line up to the field. B 3 N 3 The greater the bond polarity, the greater the dipole moment. owever, can we say that a molecule with a polar bond must also be a polar molecule? Not necessarily! CCl 4 CCl 3

7 13 14 S 4 S 6 Notice that non-polar molecules are highly symmetrical, allowing the bond dipole moments to cancel out. C. ybridization (Compounds with Single Bonds) ow does molecular shape relate to the orbitals in the valence shell? Consider methane, C 4, which uses the valence electrons in the 2s, 2p x, 2p y, and 2p z orbitals. These electrons must become unpaired prior to making bonds with. In a set of resonance structures, the overall polarity is averaged out over the resonance structures. or example, CO 3 2 and SO 4 2 are non-polar (no net dipole moment). Recall the shapes of the s and p orbitals. If we used these orbitals to bond with the atoms, we would get this structure. Yet, we know that the proper structure from VSEPR theory is tetrahedral. This structure shown is clearly incorrect! ow do we explain this?

8 Before bonding to atoms, the one 2s and three 2p atomic orbitals are mixed and rearranged to give a new set of four equivalent (same energy) hybrid atomic orbitals (sp 3 ) arranged at tetrahedral angles of Note: hybrid orbitals can also contain NB pairs, for example, in ammonia and water. These also have tetrahedral electronic arrangements and are sp 3. N 3, which has a tetrahedral electronic arrangement, contains three σ bonds. Its shape is trigonal pyramidal. 16 sp 3 hybridization is just one possibility. ive major hybridization types form the VSEPR structures, and these types are summarized in this table (details follow). Each new hybrid orbital around the C contains one electron. After each pairs up with one electron from, the orbitals contain an electron pair. The single bond formed by the direct, head-on overlap of orbitals is a sigma bond (σ). Regions of e density Atomic orbitals ybrid orbitals 2 one s, one p two sp linear Electronic arrangement 3 one s, two p three sp 2 trigonal planar 4 one s, three p four sp 3 tetrahedral 5 one s, three p, one d five sp 3 d trigonal bipyramidal 6 one s, three p, one d six sp 3 d 2 octahedral o Regions of electron density: an NB pair, a single bond, or a multiple bond each constitute one region. o Electronic arrangement may not equate to molecular shape if there is at least an NB pair present.

9 sp hybridization The combination of one s and one p results in the formation of two sp orbitals. These two hybrid orbitals are 180 apart. 2. sp 2 hybridization One s + two p = three sp 2 orbitals 120 apart. Where we have used only one of the p orbitals, there must be two p orbitals remaining. (Recall there are three p orbitals). We ll see later on that these leftover orbitals are used when there is multiple bonding. The two leftover p are 90 to each other and the sp hybrids. There is one leftover p orbital remaining, since we started with three p orbitals and used two of them for hybridization.

10 3. sp 3 hybridization One s + three p = four sp 3 orbitals apart. There are no leftover p orbitals. Also see diagram on p D. ybridization (Compounds with Multiple Bonds) ow do we assign hybridization to compounds containing multiple bonds? The same way! We still examine the number of regions of electron density (table on page 13). Consider ethene. Each carbon has an AX 3 configuration (sp 2 -hybridized). Remember to count double bonds as one region. Each C also has a remaining p orbital (see p. 15). C C Others: sp 3 d and sp 3 d 2 hybridization One s + three p + one d = five sp 3 d orbitals One s + three p + two d = six sp 3 d 2 orbitals Recall that σ bonds are formed by the direct overlap of orbitals, which result in single bonds. Where do double bonds come from? o The first bond in the double bond comes from a regular single bond caused by direct overlap (i.e. a σ bond) I 5 (sq. pyramidal)

11 o The second bond comes from the sideways overlap of the leftover p orbitals to give a pi (π) bond. This is 90 to the plane defined by the trigonal-planar σ bonds. The π bond is both above and below the plane (2 lobes). 21 ow about a triple bond? Example: acetylene o Carbon atoms are sp-hybridized, and each has two remaining p orbitals 90 to each other (see p. 14) o The first bond in the triple bond originates from the direct overlap of sp orbitals (σ bond) o Both the second and the third bonds in the triple bond are π bonds originating from the sideways overlap of the remaining p orbitals. These two π bonds are 90 apart. C C 22 Our structure therefore appears like this note that π bonds involve a sideways overlap, so bond rotation is not possible. Example: assign the hybridizations and the number of π bonds present in the molecule acetone

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