Which carbon has a covalent bond

Covalent bonds pure culture occurs when two atoms, which both have electrons, share the deficiency and combine. Examples are ...... the Halides in gas form: F.2, Cl2, Br2, etc.... many typical: O2, N2, NH3 (Ammonia), CO2, etc. such as. Si, Ge, C. (in the shape of the diamond), GaAs (Gallium arsenide) and others.The basic principle is always the same: the atoms; they are doing much better than if everyone had to cope with their electron deficiency on their own. As examples, let's look at the hydrogen and chlorine molecules, then typical ones semiconductor:
However that goes in detail in the picture above - it can be seen that the bond is undirected. After getting together with a partner, however, there is no longer any desire for further ties (or to put it correctly: there are none free valences more). We therefore describe the binding with the correct number of "binding arms". Number of bond arms = number of free valences.We can with only binding arm, i. H. if the prospective partner is fortunate enough to be only missing one electron, they clearly get crystals - we already had that.The elements that are exciting for us Si, Ge, Ga, As, N, P (not to forget: C.) but are missing 4 (or 3 or 5) electrons, and therefore they also offer suitable partners 4 (or 5 or 3) electrons to share. As soon as a suitable partner is in the vicinity - of their own kind or another suitable one (but not everyone!) - they reorganize their array of electrons in such a way that 4 Electrons preferentially in the 4Tetrahedral directions and stay in the clubs shown below, which bear the beautiful name.
With these clubs we have four clearly defined binding arms. We already had what happens when many of the atoms come together now; it is shown again above. If all atoms are of the same kind, we would have diamond (= Carbon C.), Silicon (Si) Germanium (Ge) etc .; the crystal structure is called in all cases Diamond structure. With two types of atoms, it will be that Zinc blende structure, and that concerns (in addition to the eponymous ZnS) especially GaAs, GaN, InP, GaP, ... - all pretty important to them ET&IT. Not only are all parameters unknown for the binding potential, these parameters are also still unknown. As far as we are concerned here, we can only state for further calculations:!
 What else can we learn about the properties of covalent bonds?It can come in several ways. We have to accept that because pure carbon (C.) is usually not a diamond, but a graphite. It looks like this:
Graphite structure
Each carbon atom has only coplanar bond arms, so can only form hexagonal layers, as shown in principle. Actually, you can't actually form a three-dimensional crystal with it, but the layers still hold together when placed on top of one another - via what are known as secondary bonds.Now we have a lot of questions (and answers):
  1. Why does carbon (and other atoms) do this? Simple: Because it is possible in principle.
  2. But why does carbon prefer graphite to diamond? Because the one for a bunch of carbon atoms in the formation of graphite is greater than in the formation of diamond.
  3. Why then do the diamonds that are somehow created (at high pressure and temperature) not spontaneously transform into graphite? Because you have to put energy into it first so that you can realize the energy gain resulting from the conversion. (As in real life! We (or a piece of wood) do not burn spontaneously either, although that would be energetically favorable. We would have to "ignite" us = supply energy.)
  4. Is there Si etc. also in graphite form? No. In principle it is possible, but energetically so unfavorable that it never happens.
We recognize: You can do a lot with covalent bonds (e.g. carbon chains and thus you, dear reader). Despite the enormous variety that can be drawn, a few general conclusions can be drawn:Covalently bound crystals are typically - at least at very low temperatures. Because they first have free electrons.Semiconductors are also covalently bound, but they must have a few free electrons. Yes - but only at temperatures that are high enough to use thermal energy kB.T (or light energy H) to tear a few electrons from their bonds!Covalently bound crystals are transparent to light with one energy H, which is less than the energy required to tear electrons from bonds. That's the way it is! Diamonds are transparent to visible light, Si is transparent to the near IR. Covalent bonds are strong - melting points are rather high, E-modules are rather high, the bond does not break as quickly, etc.So far so good. It is slowly becoming clear, however, that in order to understand semiconductors we have to take a closer look at the electrons in their orbitals - we need something! Here are the tasks:

© H. Föll (MaWi for ET&IT - Script)