Molecular Geometry and Bonding Theories

MOLECULAR GEOMETRY AND BONDING THEORIES

I. Molecular Geometries

A. The geometry of a molecule, along with its size, determines in large part its chemical behavior.

1. One of the most common geometries, especially in organic chemistry, is the tetrahedron.

2. A bond angle is the angle formed by the bonds of two peripheral atoms with the central atom.

B. The valenceshell electron pair repulsion (VSEPR) model is an approach to predicting geometries that considers how many electron pairs need to exist around the central atom.

C. When there are four or fewer valenceshell electron pairs around a central atom, the electronpair geometries are: linear for two regions of electrons, trigonal planar for three regions of electrons, and tetrahedral for four regions of electrons.

D. Nonbonding electrons and multiple bonds affect observed bond angles.

E. Molecules with expanded valence shells have geometries that are also predicted by VSEPR. With five regions of electrons around the central atom, the electronpair geometry is a trigonal bipyramid. With six regions of electrons, the electronpair geometry is an octahedron.

F. For molecules with no single central atom, a geometry is determined for each central atom, rather than trying to describe the overall geometry of the molecule.

Discussion Question: There are a number of tasks and devices that we can use with one hand but not the other. Have students suggest some of these things and discuss them.

II. Dipole Moments

A. A polar molecule is one in which the centers of positive and negative charge do not coincide.

1. Any diatomic molecule with a polar bond is polar.

2. A molecule’s dipole moment is a measure of the polarity.

3. The common unit of dipole moment is the debye

B. The polarity of polyatomic molecules is a function of a molecule’s geometry and the polarity of its bonds.

Discussion Question: Even though it has not yet been formally covered, ask students how polarity might affect the ability of a substance or mixture to conduct electricity.

III. Covalent Bonding and Orbital Overlap

A. Valence bond theory is a description of covalent bonding that combines Lewis’s formulation and the atomic orbital idea of wave mechanics.

B. A covalent bond forms when an atomic orbital of one atom merges, or coexists in the same space, with an atomic orbital of another atom. This merging is called orbital overlap, and the resulting covalent bonding orbital is called a valence bond orbital.

C. The valence bond orbital between two atoms is a region of high probability of finding the electron. As the nuclei approach each other, the attractive forces between the electrons and both nuclei increase, as do the repulsive forces between the nuclei. The bond length is the distance of separation at which the total energy is minimized.

D. The imaginary line that passes through both nuclei is called the internuclear axis.

E. Sigma bonds are those in which the electron density is circularly symmetrical to the internuclear axis.

F. Pi bonds are those in which the electron density is above and below the internuclear axis.

G. A node is any region around a molecule, except the outermost edges, where electron density is zero.

H. The extent of overlap tends to be greater in sigma bonds than in pi bonds.

I. In general, single bonds are sigma bonds. Double bonds consist of a sigma bond and a pi bond, and triple bonds consist of a sigma bond and two pi bonds.

IV. Hybrid Orbitals

A. Hybridization is the process of mathematically mixing two or more atomic orbitals on a single atom, giving rise to a set of blended orbitals called hybrid orbitals. The number of hybrid orbitals formed is always equal to the number of atomic orbitals used.

B. The sp hybrid orbitals are formed from the mixing of an sorbital and a porbital. The arrangement of the two sp hybrid orbitals is linear.

C. The sp2 hybrid orbitals are formed from the mixing of an sorbital and two porbitals. The arrangement of the three hybrid orbitals is trigonal planar.

D. The sp3 hybrid orbitals are formed from the mixing of an sorbital and three porbitals. The arrangement of the four hybrid orbitals is tetrahedral.

E. Hybridization can also involve dorbitals.

F. Hybridization explains the directions in which bonds point but does not address the equivalence of bonds. Promotion is a process by which an electron pair is separated into two unpaired electrons; one electron is thus promoted into a higherenergy orbital.

G. The VSEPR geometry correlates perfectly with the hybridization. The steps are given in determining the hybridization in a molecule.

H. Molecular Symmetry describes

properties of molecules due to their molecular symmetry.

Discussion Question: Choose several simple molecules and ask students to comment on the amount of symmetry each exhibits. These comments will necessarily be qualitative, unless you go into more detail on symmetry.

V. Hybrid Orbitals and Multiple Bonds

A. In many hybridizations, there are leftover unhybridized orbitals. These orbitals are the ones available for pibonding.

B. Generally, pi bonds are formed between unhybridized porbitals on the two atoms involved, with two regions of electron density on opposite sides of the internuclear axis.

C. Different pi bonds are described.

1. Localized pibonding is the normal case in which the pielectron density is entirely associated with the two atoms and their bond.

2. When double bonds exist in a molecule such that they can occur at several places, delocalized pibonding can occur. The electrons can migrate among the piregions available to them.

3. A peculiar condition exists when the alternating double and single bonds are in a cyclic arrangement with an even number of vertices. An arrangement in which six pielectrons are involved on a sixmember carbon ring is called aromaticity.

D. The valence bond theory is summarized here.

1. Every pair of bonded atoms shares one or more pairs of electrons.

2. The electrons in sigma bonds are localized in the region between two bonded atoms.

3. When atoms share more than one pair of electrons, the additional pairs are in pi bonds.

4. Electrons in pi bonds that extend over more than two atoms are said to be delocalized.

VI. Molecular Orbitals

A. Molecular orbital (MO) theory explains why covalent bonds form in terms of energy. One approach to MOs is to consider the combination of atomic orbitals from each of the atoms of the bond. Whenever two atomic orbitals interact, two MOs are formed.

B. The general features of MO theory can be illustrated by the hydrogen molecule.

1. The first MO is formed by combining the 1s atomic orbitals so that the electron density is concentrated between the nuclei. This orbital is called a bonding molecular orbital.

2. Another MO is formed by combining the 1s atomic orbitals so that the electron density is distributed away from the space between the nuclei. This orbital is called an antibonding molecular orbital.

3. The MOs formed from the 1s atomic orbitals are symmetrical with respect to the internuclear axis and thus are sigma molecular orbitals. A molecular orbital is commonly labeled with the atomic orbital that formed it.

4. The relationship among the two atomic orbitals and the two molecular orbitals can be illustrated with an energy level diagram or MO diagram.

5. The bonding MO is lower in energy than either of the contributing atomic orbitals. Electrons in this orbital are called bonding electrons.

6. The antibonding MO is higher in energy than either of the contributing atomic orbitals. Electrons in this orbital are called antibonding electrons.

C. Molecular orbital theory defines the net bonding present in a molecule, called the bond order.

VII. Molecular Orbitals for SecondPeriod Diatomic Molecules

A. MO theory can be applied to large molecules with several elements present. The homonuclear diatomic molecules of the second period are discussed here. The rules for constructing MO diagrams are given.

B. Because lowerlevel orbitals are invariably full, MO diagrams generally show only the valence shell atomic orbitals and electrons. Dilithium and diberyllium need only utilize the 2s atomic orbitals.

C. The MO diagrams for diboron through dineon utilize the 2p atomic orbitals.

D. The electron configurations, or the placements of the electrons in the MOs, relate with molecular properties of the species, particularly the magnetic properties.

1. Paramagnetism is the interaction of a substance with a magnetic field due to its unpaired electrons.

2. Diamagnetism is the interaction of a substance with a magnetic field due to its paired electrons.

E. Organic Dyes shows how MO theory

can be used to explain the absorption of light by molecules, particularly organic dyes.

Discussion Question: What is the relationship between the color of light absorbed and the color of light transmitted? Students with art experience can get this discussion going.

SAMPLE QUIZ QUESTIONS

1. What is meant by a delocalized orbital?

2. Why is the CC bond in acetylene so much shorter than that in ethylene?

3. Using VSEPR theory, why is the bond angle in ammonia 107o, whereas that in water is smaller, 104o?

4. Predict the geometries of the following molecules:

(a) SeO3; (b) XeF2.

5. Show, using a drawing, how p orbitals can interact to form a pi bond.

6. What is the hybridization used by the central atom in the following molecules: (a) OF2; (b) PCl5?

7. Which exerts a greater electrostatic repulsion on neighboring electron pairs, an unshared electron pair or a bonding electron pair? Explain.

8. Using MO theory, explain why the peroxide ion has a longer OO distance than does the oxygen molecule. Is the peroxide ion paramagnetic or diamagnetic?

9. Which hybridization is characterized by a bond angle of:

(a) 109o; (b) 120o; (c) 180o?

10. What distinguishes a sigma MO from a pi MO?

11. What is the bond order of NO, as predicted by MO theory?

12. Consider the following list of elements: He, H, C, N, P, S, O. From this list, identify the element that:

(a) Consists of diatomic molecules with triple bonds.

(b) Forms highly reactive fourmembered tetrahedra.

(c) Is found in nature as isolated atoms.

(d) Forms molecules consisting of eightmembered rings.

(e) Has an allotrope that consists of sheets of atoms, each possessing sp2 hybridization.

13. What is the shape of each of the following: (a) PCl3; (b)

ClO4; © H2Se?