Unit 3: Chemical Bonding

Section 3: The Shape of Molecules

Introduction

Molecules and Geometric Shapes

Polar Molecules

Introduction

What are the dimensions of molecules? Are they two dimensional like their portrayal on paper by Lewis structures, or are they three dimensional? What type of forces hold a molecules together to form solids and liquids? In this section, we shall look at the various structures of molecules and the forces that hold them together.

Molecules and Geometric Shapes

Molecules can be shown on paper by Lewis structures, but what are the actual dimensions of the molecule? Is it two dimensional, like a paper, or is it three dimensional like a statue? When considering the dimension of a molecule is useful to further examine the Lewis structures of molecules. For example, in the molecule CO₂, the carbon atom forms double bonds with two oxygen molecule. Since the there are two unpaired electrons on each of the oxygen molecules (each oxygen molecule on one side of the carbon atom), they tend to be as far apart as possible. To achieve this status, they form a 180 degrees bond angle with the carbon atom and each other. This molecular shape is called linear. Thus, all of the atoms in CO₂ lie in one plane.

Other planar molecular shapes that can be determined from Lewis structures is trigonal planar. BF₃ has a trigonal planar shape due to the fact is forms bonds with three florine atoms. Each florine atom has six unpaired electrons around it which tend to repel each other. However, the strongest repulsion occurs between each florine atom. As a result of this repulsion, the florine atoms try to be as far as possible. The actual (largest) angle formed between each florine atom arranged around a boron atom is 120 degrees. This forms a trigonal shape with 120 degrees being the internal angles. Because the name trigonal is short for triangle and a triangle lies on plane, thus it is a planar molecule.

We can further determine the dimensions of other Lewis structures that do not seem to lie in a plane due to constructive reasoning. Indeed, there are molecules that do not lie in a plane. For example, in CH₄, we would expect the angles between each H atom to be 90 degrees and two dimensional. However, this molecule is actual three dimensional, forming the maximum angle of 109.5 degrees between each hydrogen atom arranged around the carbon atom. This structure, backed by experimental evidence, is tetrahedral (shaped like a pyramid with equilateral triangles for faces).

The molecules we have described so far have had a central atom in which all the valence electrons are used. In molecules such as H₂O and NH₃ (ammonia), the valence electrons of the central atom are not only involved in bonding. For example, in NH₃, the nitrogen atom has three out of five valence electrons involved in bonding the hydrogen atoms. The other two valence electrons form an electron pair. As we look at the other structures, we see that this molecule meets the criteria for tetrahedral. However, instead of four atoms bonded to a central atom, an electron pair occupies the position of the fourth atom in NH₃. Due to the fact that electron to atom repulsion is less that atom to atom repulsion, the bond angles in NH₃ (107 degrees), is less that in CH₄ (109.5 degrees). This shape is called trigonal pyramidal (three dimensional). Similarly, if we apply the tetrahedral structure to the water molecule (H₂O), we find that the oxygen atom has two unpaired electron pairs that occupy the third and forth positions around it. This shape is called angular (three dimensional), and tends to form 104.5 degree bond angles (less that NH₃). The angular shape is commonly refer to as bent.

In addition to the above structures, some molecules have more that 4 positions available on the central atom. For example in PCl₅, The phosphorus atom forms five bonds with each chlorine atom. This gave rise to a new shape that has three chlorine atoms lying in one plane (trigonal planar) with one chlorine atom above and below the plane (linear). The bond angles are 120 degrees (between the three chlorine atoms in a plane) and 90 degrees (between the two atoms above or below the trigonal plane). This new shape is called trigonal bipyramidal (two tetrahedrals sharing a common side). Also, let us consider the shape of a molecule, SF₆, in which the sulfur atom forms six bonds each with florine atom. The shape of this molecule is similar to trigonal bipyramidal, but instead of a trigonal plane, four florine atoms are arranged in a regular quadrilateral. There are still two atoms, each below or above the quadrilateral plane. Hence, all the bond angles are 90 degrees. The shape is called octahedral due to fact that eight equilateral faces are present (three atoms form an equilateral face).

The means of determining the structures of the molecules described above from there Lewis structures is part of a central theory in chemistry known as the valance shell electron pair repulsion theory, or VSEPR for short (pronounced "vesper"). This theory tends to predict the shape of molecules based on orbitals, Lewis structures, and atomic properties. It useful because it is easy to comprehend, and can predict the shapes of many, but not all molecules.

Polar Molecules

Polar bonds can be used in intermolecular bonds (bonding between molecules). We know that in a polar molecule, that one side is slightly negative and one side is slightly positive. This distribution of charges are called dipoles. For example, in HCl, the chlorine atom tends to draw more of the electrons away form the hydrogen atom. This is known as a polar bond. HCl can be also be called a polar molecule due to its one polar bond. In a polar molecule, on end is slightly negative and one end is slightly positive (a dipole is present).

Some molecules have more that one polar bond. However, not all of these molecules are polar molecules. For example, in BeH₂, the Be-H are polar bonds. Hence, like the H-Cl bond, the Be-H bonds tend to form a dipole. The beryllium atom is the slightly positive end of the dipole and the hydrogen is the slightly negative end of the dipole. Since the BeH₂ molecule is linear, the dipoles pull in two opposite directions with the positive end located at the beryllium atom. As a result, the dipoles cancel out (the hydrogen atom are pulling equally in opposite directions) and the molecule is nonpolar.

The overall dipole of the BeH₂ molecule is zero. The same can be said about other nonpolar molecules such as CH₄, CO₂, and CCl₄. The shape of the molecules and the fact that each dipole is of equal force and pulling in opposite directions, leads to the determination that they are nonpolar molecules.

There are molecules that have more that one polar bonds and are polar molecules. For example, in NH₃, the nitrogen atom is drawing the negative charges away from the hydrogen atom arranged around it. The shape of the molecule tends to have the dipoles of the N-H bonds pointed in an erratic fashion. Also the unshared electron pair of the nitrogen atom repels the hydrogen atoms to an extent. Thus, the dipoles do not cancel and the overall dipole is not zero. The NH₃ molecule can be determined experimentally to be polar. Also, when NH₃ molecules meet, there is an attraction between the positive and negative ends of each molecule to the negative and positive end of another molecule, respectively. This intermolecular attraction (inter means between) of polar molecules is called a dipole-dipole force. This dipole-dipole force between molecules is different from ionic bonding and nonpolar molecular bonding.