1.4. Covalent bond

The main type of chemical bond in organic compounds is the covalent bond. Let us consider the mechanism of its formation and its main characteristics.

The simplest example of a compound with a covalent bond — the hydrogen molecule H2. A hydrogen atom consists of a positively charged nucleus and one unpaired electron occupying the 1s-orbital. When two hydrogen atoms approach each other, their electron clouds overlap:

As a result, the electron density increases in the space between the nuclei, meaning a negative charge is concentrated in that region. This negative charge attracts the positively charged nuclei, and a chemical bond forms between the atoms. A hydrogen molecule is formed , or H2. Such a bond is called covalent. The greater the overlap of the electron clouds of the atoms, the stronger the covalent bond.

The outer shell of a fluorine atom has one unpaired electron, so the valence of fluorine in compounds is one.

A single dash should be drawn from the symbol F in structural formulas:

Figure 4.2 shows the scheme of covalent bond formation in the fluorine molecule:

In the fluorine molecule, the atoms are bonded by one shared electron pair. Such a bond is called a single bond.

The outer shell of an oxygen atom has two unpaired electrons, so the valence of oxygen in compounds is two.

Two dashes should be drawn from the symbol O in structural formulas:

Figure 4.3 shows the scheme of covalent bond formation in the oxygen molecule:

In the oxygen molecule, the atoms are bonded by two electron pairs. Such a bond is called a double bond. A double bond is stronger than a single bond.

The outer shell of a nitrogen atom has three unpaired electrons, so the valence of nitrogen in compounds is three.

Three dashes should be drawn from the symbol N in structural formulas:

Figure 4.4 shows the scheme of covalent bond formation in the nitrogen molecule:

In the nitrogen molecule, the atoms are bonded by three electron pairs. Such a bond is called a triple bond. A triple bond is stronger than a single or a double bond.

Double and triple bonds share the common name multiple bonds.

In the ground state, the outer shell of a carbon atom has two unpaired electrons. However, a carbon atom readily transitions to an excited state, in which it has four unpaired electrons:

Because of this, the valence of carbon can be equal to four. In most compounds carbon is tetravalent. Therefore, four dashes should be drawn from the symbol C in structural formulas:

The most important characteristics of a covalent bond are its energy and length.

The energy that must be expended to break a chemical bond is called the bond energy. Bond energy is measured in kJ/mol. The stronger the bond, the more energy must be expended to break it, and therefore the greater the bond energy.

The energy of multiple bonds is greater than that of single bonds (Table 4.1). This is quite understandable, since breaking two or three bonds requires more energy than breaking one bond.

Atoms joined by a chemical bond are located at a certain distance from each other. The distance between the nuclei of the atoms forming the bond is called the bond length.

From the data in Table 1.4 it is clear that in the series ethane — ethylene — acetylene, as the bond order increases, the carbon-carbon bond length decreases. That is, atoms are attracted to each other more strongly when multiple bonds are formed.

In the hydrogen molecule H2, the covalent bond forms between atoms of the same element, so the shared electron pair is located symmetrically (in the middle) between the atomic nuclei. Such a bond is called nonpolar covalent.

Example 1. Covalent nonpolar bonds join together the atoms of nonmetals in simple substances

In addition, a nonpolar covalent bond forms between carbon atoms in ethane , ethylene and acetylene .

When a covalent bond forms between atoms of different chemical elements, the shared electron pair shifts toward the more electronegative atom. Such a bond is called polar covalent.

Recall that electronegativity is the ability of an atom in a chemical compound to attract the electrons of other atoms toward itself. The most electronegative element — fluorine (EN — 4,1).

Example 2. Polar covalent bond in the hydrogen fluoride molecule HF. The molecule is formed by atoms of hydrogen and fluorine. The electronegativity of fluorine is higher than that of hydrogen, so the shared electron pair is shifted toward the more electronegative fluorine atom (Fig. 4.5).

As a result, a partial negative charge arises on the fluorine atom, and a partial positive charge on the hydrogen atom. Partial, rather than full, charges arise because the shared electron pair shifts toward the fluorine atom, but a complete transfer of the electron from one atom to the other does not occur. The partial charge is denoted by the Greek letter δ (delta):

The bond between the hydrogen and fluorine atoms is polar covalent.

A polar covalent bond arises between atoms of nonmetals with different electronegativities, for example in the molecules of hydrogen chloride, water, and ammonia:

The more the electronegativities of the atoms forming the covalent bond differ from each other, the greater the partial charges on the atoms will be, and the more polar the bond will be.

Example 3. Polar covalent bonds in the molecules of hydrogen fluoride HF and hydrogen chloride HCl.

Table 4.2. Electronegativities and partial charges on atoms in the molecules HF and HCl

Substance name	Electronegativities of the elements		Partial charges on atoms
Hydrogen fluoride
Hydrogen chloride