The properties of covalent compounds are based on many factors that depend essentially on molecular structures. To begin with, the covalent bond must join your atoms and there can be no electrical charges; otherwise, one would be talking about ionic or coordination compounds.
In nature there are too many exceptions in which the dividing line between the three types of compounds becomes blurred; especially when considering macromolecules, capable of harboring both covalent and ionic regions. But generally, covalent compounds create simple, individual units or molecules.
Coast of a beach, one of the infinite examples of sources of covalent and ionic compounds. Source: Pexels.
The gases that make up the atmosphere and the breezes that hit the beach coasts are nothing more than multiple molecules that respect a constant composition. Oxygen, nitrogen, carbon dioxide, are discrete molecules with covalent bonds and are intimately involved with the life of the planet.
And on the marine side, the water molecule, OHO, is the quintessential example of a covalent compound. On the coast, it can be seen above the sands, which are a complex mixture of eroded silicon oxides. Water is liquid at room temperature, and this property will be important to keep in mind for other compounds.
Covalent bond
It was mentioned in the introduction that the gases mentioned have covalent bonds. If you take a look at their molecular structures, you will see that their bonds are double and triple: O = O, N≡N, and O = C = O. In contrast, other gases have single bonds: HH, Cl-Cl, FF and CH 4 (four CH bonds with tetrahedral geometry).
A characteristic of these bonds, and consequently of covalent compounds, is that they are directional forces; it goes from one atom to another, and its electrons, unless there is resonance, are localized. Whereas in ionic compounds, the interactions between two ions are non-directional: they attract and repel other neighboring ions.
This implies immediate consequences on the properties of covalent compounds. But, referring to its bonds, it is possible, as long as there are no ionic charges, to affirm that a compound with single, double or triple bonds is covalent; and even more so, when these are chain-type structures, found in hydrocarbons and polymers.
Some covalent compounds link into multiple bonds, as if they were chains. Source: Pexels.
If there are no ionic charges in these chains, as in the Teflon polymer, they are said to be pure covalent compounds (in a chemical and not a compositional sense).
Molecular independence
As covalent bonds are directional forces, they always end up defining a discrete structure, rather than a three-dimensional arrangement (as is the case with crystal structures and lattices). Small, medium, annular, cubic molecules, or with any other type of structure, can be expected from covalent compounds.
Small molecules, for example, include gases, water, and other compounds such as: I 2, Br 2, P 4, S 8 (with a crown-like structure), As 2, and silicon polymers and carbon.
Each of them has its own structure, independent of the links of its neighbors. To emphasize this, consider the allotrope of carbon, fullerene, C 60:
Fullerenes, one of the most interesting allotropes of coal. Source: Pixabay.
Note that it is shaped like a soccer ball. Although the balls can interact with each other, it is their covalent bonds that defined this symbolic structure; that is to say, there is not a fused network of crystalline balls, but separated (or compacted).
However, molecules in real life are not alone: they interact with each other to establish a visible gas, liquid or solid.
Intermolecular forces
The intermolecular forces that hold individual molecules together are highly dependent on their structure.
Nonpolar covalent compounds (such as gases) interact through certain types of forces (dispersion or London), while polar covalent compounds (such as water) interact by other types of forces (dipole-dipole). All these interactions have one thing in common: they are directional, just like covalent bonds.
For example, water molecules interact through hydrogen bonds, a special type of dipole-dipole forces. They are positioned in such a way that the hydrogen atoms point towards the oxygen atom of a neighboring molecule: H 2 O - H 2 O. And therefore, these interactions present a specific direction in space.
As the intermolecular forces of covalent compounds are purely directional, it means that their molecules cannot coalesce as efficiently as ionic compounds; and the result, boiling and melting points that tend to be low (T <300 ° C).
Consequently, the covalent compounds at room temperature are usually gaseous, liquid or soft solids, since their bonds can rotate, giving flexibility to the molecules.
Solubility
The solubility of the covalent compounds will depend on the solute-solvent affinity. If they are apolar, they will be soluble in apolar solvents such as dichloromethane, chloroform, toluene, and tetrahydrofuran (THF); if they are polar, they will be soluble in polar solvents, such as alcohols, water, glacial acetic acid, ammonia, etc.
However, beyond such solute-solvent affinity, there is a constant in both cases: covalent molecules do not break (with certain exceptions) their bonds or disintegrate their atoms. Salts, for example, do destroy their chemical identity when dissolving, solving their ions separately.
Conductivity
Being neutral, they do not provide an adequate medium for the migration of electrons, and therefore, they are poor conductors of electricity. However, some covalent compounds, such as hydrogen halides (HF, HCl, HBr, HI) dissociate their bond to give rise to ions (H +: F -, Cl -, Br -…) and become acids (hydracids).
They are also poor conductors of heat. This is because their intermolecular forces and the vibrations of their bonds absorb some of the supplied heat before their molecules increase in energy.
Crystals
Covalent compounds, as long as their intermolecular forces allow, can be arranged in such a way as to create a structural pattern; and thus, a covalent crystal, without ionic charges. Thus, instead of a network of ions there is a network of molecules or atoms covalently linked.
Examples of these crystals are: sugars in general, iodine, DNA, silica oxides, diamonds, salicylic acid, among others. With the exception of diamond, these covalent crystals have melting points much lower than those of ionic crystals; that is, the inorganic and organic salts.
These crystals contradict the property that covalent solids tend to be soft.
References
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