Lesson 8: Other Covalent Materials

Let's start with the simplest case, a hydrogen atom bonding to another hydrogen atom. Each has one electron and wants one more. By coming together, each can "gain" one electron from the other. Since neither atom lets go of its electron, the two atoms are bonded together by their mutual attraction for the shared pair of electrons

H₂ is an element because it contains only hydrogen atoms. H₂ is a molecule, no additional bonding is needed. There are two atoms in the molecule so it is a diatomic molecule. Hydrogen is one of several elements that form diatomic molecules.

Nitrogen atoms have five valence electrons and need three more. They are small and have a strong attraction for electrons. Thus when a nitrogen atom bonds to another nitrogen atom, each atom can attract three of the other's valence electrons and form a triple bond. Three elecrons from each atom  get lined up between the atoms.

Although that could be represented as shown in the middle diagram, triple bonds are not written this way. Generally they will be written with the three pairs of electrons arranged as shown in the bottom diagram. (Or with three parallel lines from one N to the other, shown in your workbook, ex. 12b.) So the nitrogen molecule, which can be written as N₂, has a triple bond.

Oxygen also forms a molecule called ozone that contains three atoms. If you have not already checked out the diagram shown, see if you can figure out a good Lewis diagram for O₃.

Look at the electron dot diagram for carbon. (This sequence is also shown in example 13.)

We can show the bond between this atom and another carbon in this way. But notice that this doesn't satisfy the octet rule for this first carbon atom.

To do that we'd need another carbon atom and another and another. Then each of those carbon atoms would need another and another and another and another and another and so on. This pattern, which is shown here in two dimensions, actually exists in three dimensions.

It is the diamond form of carbon, which is shown here. This model is also in the lab for you to look at when you are there. It has the same pattern of atoms as the diagram above, except that two colors are used to reperesent the carbon atoms.

However, if you start taking a look at the different dimensions of this thing, you can see that quite a different set of patterns  arise. The reason for the different colors in this model is to emphasize certain bonding patterns. You can see that there are planes or layers of atoms.

In the orientation shown to the right, you can see that the light ones are a little bit higher than the black ones. So you can see that it is a matter of position, rather than types of atoms. But that gives you a picture of what the diamond arrangement is like. Notice that each carbon atom is bonded to four other carbon atoms.

Another form of carbon is graphite and its arrangement can be envisioned by taking the puckered planes of atoms from the diamond form and making them flat. If you do that, then you end up with the graphite form of carbon. The atoms are arranged slightly differently. From the side you can see that the graphite has planes of carbon atoms that are somewhat more distant from the next plane than was the case with diamond. That is what gives graphite its slippery quality. This plane of atoms can slide on this plane of atoms because they are not quite so strongly bonded.

By changing the angle of view, you can see a hexagonal pattern. From this angle, it looks very much like the diamond. These are the two most common forms of carbon. Diamond and graphite.

Another form of the element carbon that is not as common but is making news are the Buckminsterfullerenes, also called Bucky balls, which contain clusters of carbon atoms. This particular model emphasizes the bonds. Each point where the bonds come together represents a carbon atom. These Bucky balls each contain about sixty carbon atoms, some contain 60, some 70, some 72. This is a form of carbon that has been making the news during the past several years. Carbon "nanotubes" are another form of Bucky balls, but with a tubular shape rather than a ball shape.

Take another look at carbon dioxide, CO₂, which we've talked about before. (It is also shown in example 15-a.) Because a carbon atom is smaller and has a greater pull on electrons than a silicon atom, its four electrons can be concentrated into two double bonds to the oxygen atoms.

··                ··
O : : C : : O
··                ··

Because silicon is larger and has less pull on its electrons, its electrons are spread out in four single bonds with four oxygen atoms. (Also shown in example 15-b.) This leaves each oxygen atom with the ability, actually, the necessity to bond to other atoms. Each oxygen atom shown here will bond to another silicon atom (shown in example 15-c), and each silicon atom will bond to more oxygen atoms. You can imagine that the network of covalently bonded silicon and oxygen atoms can continue indefinitely.

·
: O :
··     ··     ··
· O : Si : O ·
··     ··     ··
: O :
·

If you look carefully at this three-dimensional model of quartz, you can see that the bonding is quite extensive.

This three-dimensional model of quartz is in the lab so that you can look at it first hand when you are in the lab. At that time, you can also see that the alignment of the atoms parallels the faces of a quartz crystal.

Let me summarize a few points about covalent materials. If you have covalent bonding, you may have elements or compounds. Also, you may have either network or molecular materials. Usually it will be molecular. The only examples of network covalent bonding that you have to worry about for this course are carbon, in the form of diamond and graphite, silicon, and silicon dioxide, SiO₂, which is commonly known as quartz. Any material in this course which has just covalent bonding other than graphite, diamond, silicon, or quartz, will be a molecular material.

While I'm talking about silicon and covalent network materials, let me deal with the special case of semiconductors.

Silicon is a network marginal nonmetal. It can be represented using electron dot diagrams in this way. This is a two dimensional drawing rather than three, but it will serve for our purposes. Around the edge I've shown the pairs of electrons that would result from this atom bonding to the next atom over. Pure silicon is an intrinsic semiconductor, but impurities enhance its properties.

What if an atom from a neighboring family like phosphorus or arsenic took the place of one of the silicon atoms?

Phosphorus has five valence electrons, which is normal for it, but the bonding network is set up using 4 valence electrons from each atom. In that sense the fifth valence electron of the P is an "extra" electron, it is not "tied down" like the rest, it is somewhat free to move around, it increases the electrical conductivity of the network. Since the electron is negative, we have what is called a negative or n-type semiconductor.

Conversely, what would happen if an element like Al or Ga from the family to the left of Si on the periodic table were used to replace one of the silicon atoms?

 Ga has only 3 valence electrons. Again that is normal for Ga, but in a network that is set up using 4 valence electrons from each atom that leaves a gap. Because it represents no charge where there "should be" a negative charge, it is often referred to as a "positive hole."

The "hole" also offers a place to which neighboring electrons can move. When they do move, it is almost like the hole moved. Because the electrons can move a little, electrical conductivity is increased. Because it seems like the "positive hole" is moving, this is called a positive or p-type semiconductor.

One example is the chlorine free radical that has been implicated in the destruction of ozone in the upper atmosphere.

Another example is the hydroxyl free radical that can result when water molecules are broken apart by radiation or other reactions. This happens in our bodies often enough that we have certain enzymes dedicated to reacting with hydroxyl free radicals.

A third is nitrogen dioxide, one of the components in smog. A variety of Lewis diagrams can be drawn for this one, but they all seem to end up with one unpaired, unbonded electron.