State the type of bond formed when the combining atoms have zero electronegativity difference

By Serm Murmson

Electronegativity describes the degree to which an atom attracts electrons in a chemical bond. The difference in the electronegativity of two atoms determines their bond type. If the electronegativity difference is more than 1.7, the bond will have an ionic character. If the electronegativity difference is between 0.4 and 1.7, the bond will have a polar covalent character. Lastly, if the electronegativity difference is less than 0.4, the bond will have a nonpolar covalent character.

Unlike many other periodic trends, electronegativity does not have actual units. Instead, it is a way of combining two other periodic trends: ionization energy and electron affinity. Ionization energy is the amount of energy required to remove an electron from a neutral atom. Electron affinity is the amount of energy given off or required when a neutral atom gains an electron. Electronegativity does not have any units. However, the Pauling scale for electronegativity lists cesium as the least electronegative element, with a value of 0.79. In this scale, fluorine is the most electronegative element, with a value of 4.0.

Sodium has an electronegativity of 0.9, while chlorine has an electronegativity of 3.0. The difference between these values is 2.1, which means that sodium chloride has an ionic bond. In an ionic bond, the more electronegative element will attract an electron from the less electronegative element. In this case, chlorine becomes a Cl- ion, while sodium becomes an Na+ ion.

Hydrogen has an electronegativity of 2.0, while oxygen has an electronegativity of 3.5. The difference in electronegativities is 1.5, which means that water is a polar covalent molecule. This means that the electrons are drawn significantly towards the more electronegative element, but the atoms do not become ionized. In water, an electron from each of the hydrogen atoms is drawn towards the oxygen atom. In this case, the oxygen atom has a partial negative charge, whereas the hydrogen atoms have partial positive charges.

In hydrogen gas, two hydrogen atoms bond together. The electronegativity difference between these atoms is zero. This is a nonpolar covalent bond. In this case, the electrons of the two atoms do not favor one atom over the other. Instead, they orbit both nuclei and are said to be shared between the two atoms. Covalent bonds can exist between two different elements as well, as long as their electronegativity difference is less than 0.4.

The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). Figure \(\PageIndex{2}\) shows the relationship between electronegativity difference and bond type.

Figure \(\PageIndex{2}\): As the electronegativity difference increases between two atoms, the bond becomes more ionic.

A rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown in Figure \(\PageIndex{4}\). This table is just a general guide, however, with many exceptions. For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH3 a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI2 have a difference of 1.0, yet both of these substances form ionic compounds.

The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.

Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH–, \(\ce{NO3-}\), and \(\ce{NH4+}\), are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, KNO3, contains the K+ cation and the polyatomic \(\ce{NO3-}\) anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K+ and \(\ce{NO3-}\), as well as covalent between the nitrogen and oxygen atoms in \(\ce{NO3-}\).

Example \(\PageIndex{1}\): Electronegativity and Bond Polarity

Bond polarities play an important role in determining the structure of proteins. Using the electronegativity values in Figure \(\PageIndex{1}\), arrange the following covalent bonds—all commonly found in amino acids—in order of increasing polarity. Then designate the positive and negative atoms using the symbols δ+ and δ–:

C–H, C–N, C–O, N–H, O–H, S–H

Solution

The polarity of these bonds increases as the absolute value of the electronegativity difference increases. The atom with the δ– designation is the more electronegative of the two. Table \(\PageIndex{1}\) shows these bonds in order of increasing polarity.

Table \(\PageIndex{1}\): Bond Polarity and Electronegativity Difference

Bond	ΔEN	Polarity
C–H	0.4	\(\overset{δ−}{\ce C}−\overset{δ+}{\ce H}\)
S–H	0.4	\(\overset{δ−}{\ce S}−\overset{δ+}{\ce H}\)
C–N	0.5	\(\overset{δ+}{\ce C}−\overset{δ−}{\ce N}\)
N–H	0.9	\(\overset{δ−}{\ce N}−\overset{δ+}{\ce H}\)
C–O	1.0	\(\overset{δ+}{\ce C}−\overset{δ−}{\ce O}\)
O–H	1.4	\(\overset{δ−}{\ce O}−\overset{δ+}{\ce H}\)

Exercise \(\PageIndex{1}\)

Silicones are polymeric compounds containing, among others, the following types of covalent bonds: Si–O, Si–C, C–H, and C–C. Using the electronegativity values in Figure \(\PageIndex{3}\), arrange the bonds in order of increasing polarity and designate the positive and negative atoms using the symbols δ+ and δ–.

Answer

Bond	Electronegativity Difference	Polarity
C–C	0.0	nonpolar
C–H	0.4	\(\overset{δ−}{\ce C}−\overset{δ+}{\ce H}\)
Si–C	0.7	\(\overset{δ+}{\ce{Si}}−\overset{δ−}{\ce C}\)
Si–O	1.7	\(\overset{δ+}{\ce{Si}}−\overset{δ−}{\ce O}\)

State the type of bond is formed when the combining atoms have :

zero E.N. difference

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