Work in groups on these problems. You should try to answer the questions without referring to your textbook. If you get stuck, try asking another group for help.
“The carbon atom has four valence electrons. The spectrum shows that one of the electrons is different from the other three, and yet the four bonds of the carbon atom seem to be identical with one another....I had the idea that the electrons might occupy four equivalent tetrahedral orbitals....I worked at my desk all night, so full of excitement that I could hardly write.”
So far, we have considered, for the most part, two-dimensional or “connectivity” representations of molecules and ions. This approach has been very productive in constructing Lewis structures and discussing the chemical bond. In this workshop, we consider the three-dimensional structure of molecules. Three-dimensional structures have a different kind of importance than do two-dimensional representations. Three-dimensional representations inform us about the interactions between different molecules and interactions between different parts of the same molecule.
In this workshop we will explore one of the models used in chemistry to predict molecular structure, the Valence Shell Electron Pair Repulsion model, or VSEPR. The VSEPR model is a useful starting point for learning about the three-dimensional structure of molecules, but it does not have a firm theoretical basis. Later in this course, or perhaps in more advanced courses, you may learn about other more rigorous models. One such more complete description includes the concept of orbital hybridization, as mentioned by Professor Pauling in the unit-opening quote.
The Valence Shell Electron Pair Repulsion Model
The VSEPR model considers the interaction among the electrons within a molecule or ion as the determining factor of structure. You will use the two-dimensional Lewis diagram as the basis for the three-dimensional molecular structure.
To understand the VSEPR model, consider each “group” of electrons as an arm of electron density projecting from the central atom. Each arm repels all the other arms, so each individual arm tries to get as far away from the others as possible. In this context, we will consider a pair of electrons to be a bond if between two atoms or a "free / lone" pair if unshared with another atom. More than one pair of electrons can be shared by two atoms and there are three different types of bonds. The following list summarizes the group possibilities and bonding terms.
- A single bond has 2 shared electrons (one bonding pair). It is referred to as a σ [sigma] bond and has a bond order of one.
- A double bond has 4 shared electrons (two bonding pairs). The two bonds are distinguished from each other. One is a σ [sigma] bond and the other is referred to as a π [pi] bond. The bond order is two.
- A triple bond has 6 shared electrons (three bonding pairs). Two of the bonds are π [pi] bonds and the third is a σ [sigma] bond. The bond order is three.
- An unshared pair is referred to as a "lone" or "free" pair.
The electron-pair geometry, or arrangement of electron groups around a central atom, is based on the number of groups around that atom. Each geometry has a name which chemists use to describe the shape:
|Number of Electron Groups Around the Central Atom||Electron-Pair Geometry|
The molecular geometry is determined by the number of atoms bonded to the electron groups around the central atom. The following list is thorough, but not complete.
|Number of Electron Groups Around the Central Atom||Number of Groups of Electrons Bonded to an Atom||Molecular Geometry|
Refinement of the VSEPR Model
Lone electron pairs, which are not confined between nuclei, use more space than do bonding pairs, which are restricted to stay in the narrow space between their two nuclei. The “fat” lone pairs distort the ideal electron-pair geometries predicted by VSEPR theory. Consider the ammonia molecule as an example. It has four electron pairs, and thus VSEPR predicts a tetrahedral geometry with 109.5° bond angles. However, one of the four is a lone pair, taking up more space, and so the three bonding electron pairs are squeezed together slightly. The experimentally measured H–N–H bond angle in ammonia is, in fact, 107.3°.
A polar molecule is also known as a dipole. The distribution of charge is asymmetric in a dipole, resulting in positive and negative poles. In other words, there is a little more negative charge on one end of the molecule than on the other. The amount of “uneven-ness” in charge distribution varies among different molecules. To illustrate this concept, let’s consider CO2 and H2O, which are both three-atom molecules.
Carbon dioxide has two electron groups, both bonded, so its molecular geometry is linear. Oxygen is more electronegative than carbon, so if we considered the C=O bond alone, the bonding electrons would be asymmetrically distributed toward the oxygen atom. But the bond is not alone; there are two C=O bonds in CO2, opposite one another:
The pull toward the oxygen atom on the left exactly balances the pull toward the oxygen atom on the right. The polar bonds cancel one another. The molecule itself is nonpolar.
Water has four electron groups, with two bonded, so it has a bent molecular geometry. Oxygen is more electronegative than hydrogen, so an isolated O–H bond has its bonding electrons distributed toward the oxygen atom. Now let’s consider the overall molecule:
The pull toward the oxygen atom is not balanced by a pull in the opposite direction. In fact, both pairs of bonding electrons will be distributed toward the oxygen atom, creating a significant build-up of electrons on that end of the molecule. We indicate this with a δ–, which is used to symbolize a partial negative charge. The opposite end of the water molecule will be electron deficient, and therefore possess a partial positive charge, indicated by the symbol δ+.
A polar molecule, such as water, is said to have a dipole moment. The dipole moment of molecules can be quantitatively measured, and it is usually expressed in a unit called a debye, symbol D. The unit is named after Peter Debye, who made many important contributions to our understanding of molecular polarity. For the purpose of this discussion, we will limit our use of quantitative dipole moments to express the relative polarity of molecules. The more polar a molecule, the greater the value of its dipole moment. Nonpolar molecules have a zero dipole moment.
Water, with a dipole moment of 1.85 D, h
as one of the largest dipole moments of any molecule. This property of water molecules leads to significant effects on the macroscopic behavior of water. For example, water is able to dissolve ionic compounds because of its large dipole moment. The anomalously high boiling point of water is also a consequence of its extreme polarity.