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Molecules with odd electronic bonds are usually very reactive. These types of bonds are only stable between atoms with similar electronegativities. [14] The idea that two electrons can be divided between two atoms and serve as a link between them was first introduced in 1916 by the American chemist G.N. Lewis, who described the formation of such bonds as a result of the tendencies of certain atoms to bind to each other so that the two have the electronic structure of a corresponding noble gas atom. In a covalent bond, electrons are divided between atoms. The most stable state for an atom occurs when its valence electron layer is full, so the atoms form covalent bonds and divide their valence electrons, allowing them to reach a more stable state by filling their valence electron layer. There are different types of structures for covalent substances, including individual molecules, molecular structures, macromolecular structures, and huge covalent structures. Individual molecules have strong bonds that hold atoms together, but there are negligible tensile forces between molecules. These covalent substances are usually gases, e.g. HCl, SO2, CO2 and CH4. In molecular structures, there are low tensile forces.

These covalent substances are liquids at low boiling temperatures (such as ethanol) and low melting temperature solids (such as iodine and solid CO2). Macromolecular structures have a large number of atoms connected by covalent bonds in chains, including synthetic polymers such as polyethylene and nylon and biopolymers such as proteins and starch. Covalent lattice structures (or huge covalent structures) contain a large number of atoms connected in layers (such as graphite) or 3-dimensional structures (such as diamond and quartz). These substances have high melting and boiling points, are often brittle and tend to have high electrical resistance. Elements that have high electronegativity and the ability to form three or four bonds of electron pairs often form such large macromolecular structures. [11] Figure 3 shows the electronegativity values of the elements as proposed by one of the most famous chemists of the twentieth century: Linus Pauling (Figure 4). In general, electronegativity increases from left to right over a period of time in the periodic table and decreases by one group. Thus, the non-metals at the top right tend to have the highest electronegativities, with fluorine being the most electronegative element of all (EN = 4.0).

Metals tend to be less electronegative, and Group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this number because these atoms usually do not share electrons with other atoms because they have a complete valence layer. (Although noble gas compounds such as XeO2 exist, they can only be formed under extreme conditions and therefore do not exactly match the general pattern of electronegativity.) Compounds consisting of covalent bonds usually have differences in physical properties (e.g. B, water solubility, conductivity, boiling point and melting point) relative to ionic compounds. In three-centric two-electron bonds (“3c-2e”), three atoms divide two electrons in the bond. This type of bonding occurs in boron hydrides such as diborane (B2H6), which are often described as low in electrons because there are not enough valence electrons to form localized bonds (2-center-2 electrons) that connect all atoms. However, the more modern description with 3c-2e bonds provides enough binding sorbitals to bind all atoms so that molecules can instead be classified as accurate by electrons. Fig. 2.

(a) Binding and (b) Anti-binding molecular orbitals of the H2 molecule, (c) Schematic drawing of the structure of the most important molecular orbitals from atomic orbitals and (d), (e) Examples of molecular orbitals (bond: σ, π and anti-bond σ*, π*) The ligand is any neutral atom, ion or molecule capable of giving a pair of electrons and bound to the ion or central metal atom by secondary valence. Covalent coordinate bonds are covalent bonds in which the two bonding electrons are brought by one of the bonding partners. Figure 2 distinguishes the covalent bonds from the covalent bond in NH3BF3. While the three BF-covalent bonds are formed due to the sharing of electron pairs resulting from the contributions of boron and fluorine atoms, an NB bond is formed due to the donation of a single pair of nitrogen electrons into the boron empty orbitals. The bond of covalent coordinates is represented by an arrow whose head points in the direction of the donation of a pair of electrons, as shown in Figure 2. Figure 2. a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small black dots indicate the position of the hydrogen and chlorine nuclei in the molecule.

(b) The symbols Δ+ and Δ– indicate the polarity of the H-Cl bond. .

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