Chapter 9
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The document discusses molecular geometry and molecular orbitals, emphasizing that Lewis structures indicate atomic connectivity but not molecular shape, which is determined by bond angles. It introduces Valence Shell Electron Pair Repulsion (VSEPR) theory to predict molecular shapes based on electron pair repulsion and outlines hybridization concepts necessary for forming bonds. Additionally, it explains how molecular polarity arises from the geometry of molecules and the nature of bonding interactions, including sigma and pi bonds.
Chapter 9
- 1. Chapter 9Chapter 9 Molecular Geometry, MolecularMolecular Geometry, Molecular OrbitalsOrbitals
- 2. • Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which. However, do not show their overall shape) • A molecule’s shape is determined by its bond angles. • Consider CCl4: experimentally we find all Cl-C-Cl bond angles are 109.5°. • Therefore, the molecule cannot be planar. • All Cl atoms are located at the vertices of a tetrahedron with the C at its center. Molecular ShapesMolecular Shapes Molecular Shape of CCl4Molecular Shape of CCl4
- 3. • In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts which ever 3D geometry minimized this repulsion. • We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory. • There are simple shapes for AB2 and AB3 molecules. VSEPR TheoryVSEPR Theory
- 4. • 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. Naming Molecular Geometry
- 7. To determine the electron pair geometry: • draw the Lewis structure, • count the total number of electron pairs around the central atom, • arrange the electron pairs in one of the above geometries to minimize e− -e− repulsion, and count multiple bonds as one bonding pair. VSEPR ModelVSEPR Model
- 11. • We determine the electron pair geometry only looking at electrons. • We name the molecular geometry by the positions of atoms. • We ignore lone pairs in the molecular geometry. • All the atoms that obey the octet rule have tetrahedral electron pair geometries. The Effect of Nonbonding Electrons and multiple bonds on angles 104.5O 107O N HH H C H HH H 109.5O O HH By experiment, the H-X-H bond angle decreases on moving from C to N to O: Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. Therefore, the bond angle decreases as the number of lone pairs increase. •Similarly, electrons in multiple bonds repel more than electrons in single bonds. C O Cl Cl 111.4o 124.3o
- 12. • Atoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (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. • To minimize e--e- repulsion, lone pairs are always placed in equatorial positions. Molecules with Expanded Valence Shells
- 13. • In acetic acid, CH3COOH, there are three central atoms. • We assign the geometry about each central atom separately. Shapes of Larger Molecules
- 14. • When there is a difference in electronegativity between two atoms, then the bond between them is polar. • It is possible for a molecule to contain polar bonds, but not be polar. • For example, the bond dipoles in CO2 cancel each other because CO2 is linear. Molecular Shape and Molecular PolarityMolecular Shape and Molecular Polarity
- 15. Molecular Polarity of HMolecular Polarity of H22OO • In water, the molecule is not linear and the bond dipoles do not cancel each other. • Therefore, water is a polar molecule. • The overall polarity of a molecule depends on its molecular geometry
- 17. • Lewis structures and VSEPR do not explain why a bond forms. • How do we account for shape in terms of quantum mechanics? • What are the orbitals that are involved in bonding? • We use Valence Bond Theory: • Bonds form when orbitals on atoms overlap. • There are two electrons of opposite spin in the orbital overlap. Covalent Bonding and Orbital OverlapCovalent Bonding and Orbital Overlap
- 18. Covalent Bonding and Orbital Overlap MechanismCovalent Bonding and Orbital Overlap Mechanism • As two nuclei approach each other their atomic orbitals overlap. • As the amount of overlap increases, the energy of the interaction decreases. • At some distance the minimum energy is reached. • The minimum energy corresponds to the bonding distance (or bond length). • As the two atoms get closer, their nuclei begin to repel and the energy increases. • At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).
- 19. • Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding. • Hybridization is determined by the electron domain geometry. sp Hybrid Orbitals • Consider the BeF2 molecule (experimentally known to exist): • Be has a 1s2 2s2 electron configuration. • There is no unpaired electron available for bonding. • We conclude that the atomic orbitals are not adequate to describe orbitals in molecules. • We know that the F-Be-F bond angle is 180° (VSEPR theory). • We also know that one electron from Be is shared with each one of the unpaired electrons from F. We assume that the Be orbitals in the Be-F bond are 180° apart. We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding. BUT the geometry is still not explained Hybrid OrbitalsHybrid Orbitals
- 20. sp Hybrid Orbitals • We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital.. • The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital. • The lobes of sp hybrid orbitals are 180º apart. • Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.
- 21. • Important: when we mix n atomic orbitals we must get n hybrid orbitals. • sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.) • The large lobes of sp2 hybrids lie in a trigonal plane. • All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom. sp2 Hybrid Orbitals
- 23. • sp3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes. • Each lobe points towards the vertex of a tetrahedron. • The angle between the large lobs is 109.5°. • All molecules with tetrahedral electron pair geometries are sp3 hybridized. sp3 Hybrid Orbitals
- 25. • Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron domain geometries must involve d-orbitals. • Trigonal bipyramidal electron domain geometries require sp3 d hybridization. • Octahedral electron domain geometries require sp3 d2 hybridization. • Note the electron domain geometry from VSEPR theory determines the hybridization. Hybridization Involving d Orbitals
- 28. ∀ σ-Bonds: electron density lies on the axis between the nuclei. • All single bonds are σ-bonds. ∀ π-Bonds: electron density lies above and below the plane of the nuclei. • A double bond consists of one σ-bond and one π-bond. • A triple bond has one σ-bond and two π-bonds. • Often, the p-orbitals involved in π-bonding come from unhybridized orbitals. Multiple BondsMultiple Bonds
- 29. • Ethylene, C2H4, has: • one σ- and one π-bond; • both C atoms sp2 hybridized; • both C atoms with trigonal planar electron pair and molecular geometries. Multiple BondsMultiple Bonds EthyleneEthylene
- 30. Multiple BondsMultiple Bonds • Consider acetylene, C2H2 – the electron pair geometry of each C is linear; – therefore, the C atoms are sp hybridized; – the sp hybrid orbitals form the C-C and C-H σ-bonds; – there are two unhybridized p-orbitals; – both unhybridized p-orbitals form the two π-bonds; – one π-bond is above and below the plane of the nuclei; – one π-bond is in front and behind the plane of the nuclei. • When triple bonds form (e.g. N2) one π-bond is always above and below and the other is in front and behind the plane of the nuclei. AcetyleneAcetylene
- 31. • Every two atoms share at least 2 electrons. • Two electrons between atoms on the same axis as the nuclei are σ bonds. ∀ σ-Bonds are always localized. • If two atoms share more than one pair of electrons, the second and third pair form π-bonds. • When resonance structures are possible, delocalization is also possible. Multiple Bonds-Multiple Bonds- General Conclusions