Covalent bonds form between atoms with similar electronegativity values. Electrons are shared among atoms in a covalent bond.
When two atoms with similar electronegativities are close to each other, the atoms pull on each other's electrons with equal force. However, neither atom can pull away the other's electrons, and ion formation does not occur. The forces acting on the atoms are similar to those of ionic bonding. The protons of the two atoms repulse each other, and so do the electrons. At the same time, the protons of each atom attract the electrons of the other. The point where these two forces balance each other represents the location of minimum energy.
When attractive and repulsive forces between covalently bonded atoms are at equilibrium, the potential energy is minimized.
In this case, at the minimum energy configuration, the two atoms share electrons, forming a covalent bond. A covalent bond is a chemical bond that forms when valence electrons are shared between atoms. Covalently bonded atoms form molecules. Many substances, including water (H2O), carbon dioxide (CO2), oxygen gas (O2), and nitrogen gas (N2), are molecular.
In a covalent bond, the nuclei of both atoms exert a force on the shared electrons. When the shared electrons are to the side of the nuclei, they are attracted toward both nuclei. When the shared electrons are between nuclei, the nuclei apply opposing forces on the electron.
Forces Acting on an Electron in a Covalent Bond
When the shared electrons in a covalent bond are to the side of the nuclei, the attractive forces of the nuclei originate from the same direction relative to the electron. When the shared electrons are between nuclei, the nuclei apply opposing forces on the electrons. The strength of the nuclear force on the electrons therefore varies with the position of the electrons.
The direction of the nuclear forces depends on the positions of the electrons. So, shared electrons in a covalent bond are more likely to be found in regions between the two nuclei.
Electron Distribution in a Nonpolar Diatomic Molecule
The + signs represent the positively charged nuclei of two covalently bonded atoms. The color around the + signs represents electrons around the nuclei. In a covalent bond, electrons are more likely to be located in regions in between the two atoms at any given moment, as shown by the high density of color.
Two concepts play an important role in defining covalent bonds: bond length and bond strength. Bond length is the distance between the nuclei of the atoms forming the bond, usually expressed in picometers (pm) or angstroms (Å). Bond length is affected by the size of the atoms. Hydrogen atoms are small because they have a single proton. Elemental hydrogen (H2), therefore, has a short bond length of 74 pm, where 1 pm is equal to 10–12 meter (m). Chlorine atoms are larger because they each have 17 protons. Elemental chlorine (Cl2), therefore, has a larger bond length of 199 pm.
Bond strength indicates how much energy is required to break a bond. Hydrogen-hydrogen bonds in elemental hydrogen, for example, require 432 kilojoules per mole (kJ/mol) of energy to break. Chlorine-chlorine bonds in elemental chlorine are weaker, at 240 kJ/mol.
The molecule that a bond is a part of affects the bond strength and the bond length to a degree. For example, in methane (CH4) the length of each C−H bond is 108.7 pm. In ethane (C2H6) the length of each C−H bond is 109.4 pm. This difference in bond length and bond strength is relatively small.
A single bond is a covalent bond in which one pair of electrons is shared between two atoms. In many cases atoms can share more than one pair of electrons. A double bond is a covalent bond in which two pairs of electrons are shared between two atoms. A triple bond is a covalent bond in which three pairs of electrons are shared between two atoms. The number of shared electron pairs is called the bond order. First-order bonds share one pair of electrons, second-order bonds share two pairs, and third-order bonds share three pairs of electrons. Higher-order bonds between the same two elements tend to have shorter bond lengths and stronger bonds than single bonds between the same two elements.
Bond Strength and Bond Length of Common Bonds
Bond
Bond Strength (kJ/mol)
Bond Length (pm)
H−H
432
74
H−C
411
109
H−F
565
92
H−Cl
428
127
H−Br
362
141
H−I
295
161
C−C
346
154
C=C
602
134
C≡H
835
120
C−O
358
143
C=O
799
120
O=O
494
121
N−N
167
145
N=N
418
125
N≡N
942
110
Various factors affect bond strength and bond length, including molecular structure and the sizes of the atoms. In general, higher bond energies are associated with shorter bond lengths because the forces between the atoms are stronger.
Some trends may be observed when bond lengths and bond strengths are studied. When elements from the same group are involved in similar bonds, such as H−F, H−Cl, H−Br, and H−I, the bond length tends to increase and the bond strength tends to decrease with increasing atomic size. Double and triple bonds of the same two elements tend to have shorter bond lengths and higher bond energies compared to single bonds between the same two elements.
