Molecular Geometry and Bonding Theories. Molecular Shapes. Molecular Shapes. Chapter 9 Part 2 November 16 th, 2004

Transcription

1 Molecular Geometry and Bonding Theories Chapter 9 Part 2 November 16 th, Molecular Shapes When considering the geometry about the central atom, we consider all electrons (lone pairs and bonding pairs). When naming the molecular geometry, we focus only on the positions of the atoms. 9 Molecular Shapes To determine the shape of a molecule, need to distinguish between lone pairs (or non-bonding pairs, those not in a bond) of electrons and bonding pairs (those found between two atoms). The electron domain geometry is defined by the positions in 3D space of ALL electron pairs (bonding + non-bonding). This becomes especially important when the central atom has lone pairs. The electrons adopt an arrangement in space to minimize e - -e - repulsion. 10 1

2 11 12 Lone pairs on Central Atom Consider NH 3 : The Lewis structure of NH 3 is: The molecular geometry of NH 3 is described as trigonal pyramidal with the N atom at the apex and the 3 H atoms forming the base of the triangle. The electron-pair geometry is described as tetrahedral. 13 2

3 14 Lone Pairs & Bond Angles The electron pair geometry is decided only by looking at the positions of the electrons. The molecular geometry is decided only by the positions of atoms, lone pairs are ignored in the molecular geometry. Lone pairs are considered to be fat, while bonding pairs are considered skinny. The increased volume of lone pairs makes bond pairs squeeze together, reducing bond angles. 15 Lone Pairs & Bond Angles 16 3

4 Lone Pairs & Bond Angles Experimentally, the H-X-H bond angle decreases on moving from C to N to O: H H C H N O H109.5 H H H 107 O H H O O Since the electrons in a bonding pair are attracted by two nuclei they do not repel as much as lone pairs. The bond angle decreases as the number of lone pairs on the central atom increases. Typically the relative strength of repulsion is: 17 Multiple Bonds & Bond Angles Similarly, electrons in multiple bonds repel more than electrons in single bonds. A multiple bond is counted as one electron domain. The Lewis structure of COCl 2 (carbonyl chloride) indicates the molecular and the electron pair geometry to be trigonal-planar. Cl o C O Cl o 18 Molecules with Expanded Valence Shells Atoms that have expanded octets have AB 5 (trigonal bipyramidal) or AB 6 (octahedral) electron pair geometries. For trigonal bipyramidal structures there is a plane containing three electrons pairs. The fourth and fifth electron pairs are located above and below this plane. For octahedral structures, there is a plane containing four electron pairs. Similarly, the fifth and sixth electron pairs are located above and below this plane. 19 4

5 20 21 Molecules with Expanded Valence Shells To minimize e - -e - repulsion, lone pairs are always placed in equatorial positions. In an octahedron, all positions are considered to be equivalent and there are no axial or equatorial positions. If the molecule has any lone pairs they are placed above or below the plane. 22 5

6 Shapes of Large Molecules In acetic acid, CH 3 COOH, there are three central atoms. We assign the geometry about each central atom separately. 23 Determining Molecular Shapes Draw the Lewis structure. Determine the total number of bonding pairs (count any multiple bonds as 1 pair) around the central atom. Pick the appropriate electron pair geometry and then choose the molecular shape that matches the total # of single-bond pairs and lone pairs. Predict the bond angles, remembering that the lone pairs occupy more volume than bonding pairs. 24 Molecular Shape & Polarity Difference in electronegativity between two atoms makes the bond between them polar. It is possible for a molecule to contain polar bonds, but not be polar. For example, the bond dipoles in CO 2 cancel each other because CO 2 is linear. 25 6

7 Molecular Shape & Polarity In water, the molecule is not linear and the bond dipoles do not cancel each other. Therefore, water is a polar molecule. 26 Molecular Shape and Molecular Polarity 27 Questions not answered by VSEPR Why does a bond form? How do we account for molecular shape in terms of quantum mechanics, meaning in terms of orbital shapes? What are the orbitals that are involved in bonding? Valence Bond Theory proposed by Linus Pauling provides these answers. Provides a qualitative visual picture of molecular structure and bonding. 28 7

8 Valence Bond (VB) Theory Orbitals overlap to form a bond between atoms. Two electrons of opposite spins can be accommodated in the overlapping orbitals. Usually one electron is supplied by each of the 2 bonded atoms. Overlapping orbitals result in a high probability of the 2 electrons being located in the region that is influenced by both nuclei. Both electrons are simultaneously attracted by both nuclei Orbital Overlap Orbital overlap produces a single bond that is known as a sigma (σ) bond. The electron density of the sigma bond is greatest along its axis. The 2s and 2p electrons that are not involved in the bond are the lone pairs on the F atoms. 31 8