Covalent bond
A covalent bond is produced in two non-metallic atoms when they unite and share one or more electrons of the last level, example of throwing two and one on the dice the first result of the logical sequence but not the whole as a double of ones (valence electrons) (except hydrogen which reaches stability when it has 2 electrons) to thus reach the octet rule. The difference in electronegativity between the atoms is not large enough for an ionic bond to occur. For a covalent bond to be generated, the difference in electronegativity between atoms must be less than 1.7.
In this way, the two atoms share one or more electron pairs in a new type of orbital, called a molecular orbital. Covalent bonds occur between atoms of the same non-metal element, between different non-metals, and between a non-metal and hydrogen.
When atoms other than non-metals unite in an ionic form, one of them will be more electronegative than the other, so it will tend to attract the electron cloud of the bond towards its nucleus, generating an electric dipole. This polarization allows that molecules of the same compound are attracted to each other by electrostatic forces of different strengths.
On the contrary, when atoms of the same non-metallic element join covalently, their electronegativity difference is zero and dipoles are not created. Molecules have practically no attraction for each other.
In summary, in an ionic bond, the transfer of electrons from one atom to another occurs and in the covalent bond, the bonding electrons are shared by both atoms. In the covalent bond, the two non-metallic atoms share one or more electrons, that is, they join through their electrons in the last orbital, which depends on the atomic number in question. Between the two atoms one, two or three pairs of electrons can be shared, which will give rise to the formation of a single, double or triple bond respectively. In the Lewis structure, these bonds can be represented by a small line between the atoms.
History
The term "covalence" in relation to the union was used for the first time in 1919 by Irving Langmuir in an article in the Journal of the American Chemical Society entitled «The Arrangement of Electrons in Atoms and Molecules» (The distribution of electrons in atoms and molecules). In it, Langmuir wrote: "we will designate by the term covalence the number of pairs of electrons that a given atom shares with its neighbors".
The idea of covalent bonding can be traced back several years to Gilbert N. Lewis, who in 1916 described the exchange of electron pairs between atoms. He introduced Lewis notation or electron dot notation or Lewis dot structure, in which the valence of electrons (those in the outer shell) are represented as dots around the atomic symbols. The pairs of electrons located between atoms represent covalent bonds. Multiple pairs represent multiple bonds, such as double bonds and triple bonds. An alternate form of representation, not shown here, has the bonding pairs of electrons represented as solid lines.
Lewis proposed that an atom forms enough covalent bonds to form a complete outer (or closed) electronic shell. In the methane diagram shown here, the carbon atom has a valence of four and is therefore surrounded by eight electrons (the octet rule), four from carbon itself and four from the hydrogens attached to it. Each hydrogen has a valence of one and is surrounded by two electrons (a duet rule), its own electron plus one from carbon. The number of electrons correspond to complete shells in the quantum theory of the atom; the outer shell of a carbon atom is the n=2 shell, holding eight electrons, while the outer (and only) shell of a hydrogen atom is the n=1 shell, holding two or more elements.
While the idea of shared pairs of electrons provides an effective qualitative picture of covalent bonding, quantum mechanics is necessary to understand the nature of these bonds and to predict the structures and properties of simple molecules. Walter Heitler and Fritz London gave the first successful explanation of a chemical bond using quantum mechanics, specifically molecular hydrogen, in 1927. Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is a good match between the atomic orbitals of the participating atoms.
It is known that these atomic orbitals have specific angular relationships to each other, and therefore the valence bond model can successfully predict the bond angles observed in simple molecules.
However, the covalent bond theory or also known as the idea of electron sharing, which is based on the cubic atom, faced several conceptual difficulties, since this theory had competition from the ionic bond model. Despite this rivalry of these two theories, the covalent bond theory was accepted until 1920. M.Niaz and M.A.Rodríguez mention in their text History and philosophy of sciences: the need for its incorporation into university science textbooks that Lewis recognizes that the cubic structure cannot represent the triple bond and suggests replacing it for the tetrahedral atom. Lewis assumed for many years that if the electrons in the atom are magnetically paired, it is easy to understand how two unpaired electrons in different atoms can magnetically couple and form the nonpolar bond.
Types of covalent substances
There are two types of covalent substances:
Molecular covalent substances: Covalent bonds form molecules that have the following properties:
- Low fusion and boiling temperatures.
- In normal pressure and temperature conditions (25 °C approx.) can be solid, liquid or gaseous.
- They're soft in solid state.
- They are insulation of the electric current and heat.
- Solubility: Polar molecules are soluble in polar solvents and apolars are soluble in apolar solvents (like dissolves the like).
- Examples: carbon dioxide, benzene, oxygen, nitrogen.
Lattice covalent substances or lattices: In addition, covalent substances form crystalline lattices of an indefinite number of atoms, similar to ionic compounds, which have these properties:
- High fusion and boiling temperatures.
- They are solid in normal conditions.
- They're very hard substances.
- They are insulating (except graphite).
- They're insoluble.
- Examples: quartz, diamond.
Covalent Bond Definition
Consider hydrogen atoms, as they approach each other, the forces that attract each electron to the nucleus of the other atom become noticeable, until said attractive forces are offset by the repulsion that the electrons feel each other. At that point, the molecule presents the most stable configuration.
What has happened is that the orbitals of both electrons have overlapped, so that it is now impossible to tell which atom each of the electrons belongs to.
According to the chemists S. Seese and G. William Daub, in the hydrogen molecule, as in all covalent substances, four aspects must be taken:
First: the properties of individual uncombined atoms are very different from the properties of molecules. Therefore, when writing the chemical formula of hydrogen, it must be written as H2, because it is a diatomic molecule.
Second: the two positive nuclei attract the two electrons in order to produce a more stable molecule than that of the separated atoms, consequently a bond is formed and with it a more stable molecule results stable than before. Due to the attraction exerted by the nuclei on the two electrons, the repulsion between them is balanced and therefore the probability of finding electrons somewhere between the nuclei is greater.
Third: the distance between the nuclei is such that the 1s orbitals have the maximum overlap. In the case of the hydrogen molecule, the distance between the nuclei is approximately 0.74 Å. Otherwise, the distance between two covalently bonded atoms is called the bond length.
Fourth: to "break" covalent bonds in 1.0 g of hydrogen gas and to form hydrogen atoms 52.0 kcal are needed.
However, when the atoms are different, the shared electrons will not be equally attracted, so that they will tend to approach the most electronegative atom, that is, the one that has a greater appetite for electrons. This phenomenon is called polarity (atoms with higher electronegativity gain a more negative polarity, drawing shared electrons towards their nucleus), and it results in a shift of charges within the molecule.
It could be said that the most electronegative atom does not like to share its electrons with the other atoms, and in the most extreme case, it will want the electron to be ceded to it without conditions, thus forming an ionic bond. Hence, it is said that polar covalent bonds have, to some extent, ionic character.
When the difference has a value from 0 to 1.7, the covalent character will predominate, as is the case of the C-H bond. However, according to chemist Raymond Chang, this electronegativity difference between atoms must be 2.0 or greater for the bond to be considered ionic (Chang, 371).
Depending on the difference in electronegativity, the covalent bond can be classified as polar covalent and pure or nonpolar covalent. If the difference in electronegativity is between 0.4 and 1.7 it is a polar covalent bond, and if it is less than 0.4 it is nonpolar covalent.
When the difference in electronegativities is null (two equal atoms), the bond formed will be pure covalent; for a difference in electronegativities of 1.9, the ionic character already reaches 35%, and for a difference of 3, it will be 49.5%.
Between oxygen or fluorine and the elements of groups 1 and 2, this difference will be maximum and its ionic character too.
Polar covalent bond
In a polar covalent bond, electrons are shared unequally between atoms and spend more time near one atom than the other. Due to the uneven distribution of electrons between atoms of different elements, slightly charged charges appear. positive (δ+) and slightly negative (δ–) in different parts of the molecule.
In a water molecule, the bond that joins oxygen to each hydrogen is a polar bond. Oxygen is a much more electronegative atom than hydrogen, so the oxygen in water carries a partially negative charge (has a high electron density), while the hydrogens carry partially positive charges (have a low electron density).
In general, the relative electronegativity of the two atoms in a bond, that is, their tendency to hoard the shared electrons, will determine whether the bond is polar or nonpolar. As long as one element is significantly more electronegative than another, the bond between them will be polar; this means that one end will have a slightly positive charge and the other a slightly negative charge.
That is, it consists of the formation of a bond between atoms of different elements, and the difference in electronegativity must be greater than 0.4. In this bond, the electrons are fundamentally attracted by the nucleus of the most electronegative atom, generating molecules whose electronic cloud will present an area with a higher negative charge density and another with a higher positive charge density (dipole).
Nonpolar covalent bond
Nonpolar covalent bonds are formed between two atoms of the same element or between atoms of different elements that share electrons more or less equally. For example, molecular oxygen is nonpolar because electrons are shared equally between atoms. two oxygen atoms.
Another example of a nonpolar covalent bond can be found in methane. Carbon has four electrons in its outer shell and requires four more to become a stable octet. It achieves them by sharing electrons with four hydrogen atoms, each of which provides it with one electron. Similarly, hydrogen atoms need one extra electron each to fill their outermost shell, which they receive in the form of shared electrons from carbon. Although carbon and hydrogen do not have exactly the same electronegativity, they are quite similar, so carbon-hydrogen bonds are considered nonpolar.
| Non-polar substances | Polar substances |
|---|---|
| Etano | Methanol |
| Tolueno | Fenol |
| isobutano | Acetona |
| Pentano | Propane acid |
Different types of covalent bonds
- Simple link: It is a shared electronic pair formed by an electron belonging to the last level of energy of each atom and is represented with a line. Examples: H-H, Cl-Cl
- Double link: Formed by two shared electronic pairs, i.e. by two electrons belonging to the last level of energy of each atom and is represented with two parallel lines. Example: O=O
- Triple link: Formed by three shared electronic pairs, i.e. by three electrons belonging to the last level of energy of each atom and is represented with three parallel lines. Example: N≡N
- Dantive covalent link or coordination: It is an electronic pair shared by two atoms but both electrons are provided by the same atom. It is usually represented with an arrow (→).
An example of a chemical species that possesses a coordinate bond is the ammonium ion (NH41+). The ammonium ion is made up of a proton and ammonia.
Compounds in which a coordinate bond is found are known as coordination compounds.
Coordination compounds, also called complexes, are usually attached to several surrounding anions known as ligands.