# Polar Covalent Bond: Definition, Electronegativity & Examples

A polar bond is a covalent bond between two atoms where the electrons forming the bond are not equally distributed. As a result, the molecule has a slight electrical dipole moment, with one end being slightly positive and the other slightly negative. As the electric dipoles have less charge than a full unit charge, they are referred to as partial charges, denoted by delta plus $$\mathbf{(δ^+)}$$ and delta minus $$\mathbf{(δ^-)}$$. Molecules with polar covalent bonds are able to interact with molecules with dipoles because the positive and negative charges are separated in the bond. Intermolecular forces are produced between molecules as a result of dipole-dipole interactions.

The polar bond is the dividing line between pure covalent bonds and pure ionic bonds. Atoms in pure covalent bonds (nonpolar covalent bonds) share electron pairs equally. Nonpolar bonding is technically only possible when the atoms are identical (e.g., $$\mathbf{H_2\ gas}$$), but chemists consider any nonpolar covalent bond between atoms with a difference in electronegativity of less than $$0.4$$ to be a nonpolar covalent bond. Nonpolar molecules include carbon dioxide $$\mathbf{(CO_2)}$$ and methane $$\mathbf{(CH_4)}$$.

## Polar Covalent Bond

When atoms with different electronegativities share electrons in a covalent bond, they form a polar covalent bond. Consider the hydrogen chloride molecule $$\mathbf{(HCL)}$$. The inert gas electron configuration requires an additional electron for each atom in $$\mathbf{HCL}$$. Although chlorine has a higher electronegativity than hydrogen, its attraction for electrons is not enough to remove an electron from hydrogen.

Thus, the electrons in hydrogen chloride are shared unequally in a polar covalent bond. The molecule is represented by the conventional Lewis structure, even though the shared electron pair is associated more strongly with chlorine than with hydrogen.

An unequal sharing of the bonding pair results in a partial negative charge on the chlorine atom and a partial positive charge on the hydrogen atom. The symbol $$\mathbf{(δ)}$$ (Greek lowercase delta) denotes these fractional charges.

A hydrogen chloride molecule has a dipole (two poles), which is made up of opposite charges separated from each other. One end of the dipole is marked with an arrow and a cross. The cross is near the partially positive end of the molecule, and the arrowhead is near the partially negative end.

The bonds between carbon atoms are nonpolar. Hydrogen and carbon have similar electronegativity values, so the hydrogen-carbon bond is not normally considered a polar covalent bond. Therefore, ethane, ethylene, and acetylene have nonpolar covalent bonds, and the compounds are nonpolar.

Polar bonds exist between carbon and other elements such as oxygen and nitrogen. The electronegativities of the bonded atoms determine the polarity of a bond.

A large difference between the electronegativities of the bonded atoms increases the polarity of bonds. It is easy to predict the polarity of organic molecules based on the common bonds they contain. Nonmetallic elements tend to be more electronegative than carbon atoms. Thus, when carbon atoms are bonded to nonmetal atoms, they have a partial positive charge.

Hydrogen is less electronegative than common nonmetals. Thus, when a hydrogen atom is bonded to common nonmetals, the resulting polar bond has a partial positive charge on the hydrogen atom.

The dipole moment, $$\mathbf{(D)}$$, is a measure of the magnitude of the polarity of a bond. The dipole moment of a bond is fairly constant from compound to compound. Polar bonds occur when carbon forms multiple bonds with other elements. The carbon-oxygen double bond in formaldehyde (methane) and the carbon-nitrogen triple bond in acetonitrile (cyanomethane) are both polar.

### But Aren’t Ionic Bonds Polar?

The electrons in ionic bonds are essentially donated by one atom to the other (e.g., $$\mathbf{NaCl}$$). Atoms form ionic bonds when their electronegativity difference is greater than $$\mathbf{1.7}$$. Ionic bonds are technically polar bonds, so the terminology can be confusing.

A polar bond refers to a type of covalent bond in which electrons are not equally shared and electronegativity values are slightly different. Atoms with electronegativity differences between $$\mathbf{0.4}$$ and $$\mathbf{1.7}$$ form polar covalent bonds.

### How Electronegativity Affects Polar Covalent Bond

Since the tendency of an element to gain or lose electrons is so important in determining its chemistry, many methods have been developed to describe this tendency quantitatively. The most important method uses a measurement called electronegativity (represented by the $$\mathbf{Greek\ letter\ chi,\ χ,\ pronounced\ “ky”\ as\ in\ “sky”}$$), which is defined as the relative ability of an atom to attract electrons to itself in a chemical compound.

The elements with high electronegativities tend to acquire electrons in chemical reactions and are situated in the upper right corner of the periodic table. Low electronegativities tend to lose electrons in chemical reactions, which is why they are found toward the bottom of the periodic table.

The electronegativity of an atom is not a simple, fixed property that can be directly measured in a single experiment. It makes sense that an atom’s electronegativity should be affected by its chemical environment since the properties of an atom are affected by its neighboring atoms.

Even so, when different methods for measuring the electronegativity of an atom are compared, they tend to assign similar relative values to a given element. Linus Pauling, one of the most famous chemists of the twentieth century, proposed the electronegativity values of the elements in the figure below.

A value of $$\mathbf{4.0}$$ is arbitrarily assigned to the most electronegative element, fluorine, and the other electronegativities are scaled based on this value. Generally, electronegativity increases from left to right across a period in the periodic table and decreases down a group.

Consequently, the nonmetals in the upper right tend to have the highest electronegativities, with fluorine being the most electronegative element ($$\mathbf{EN = 4.0}$$ as previously stated). The elements most important for organic chemistry, carbon, nitrogen, and oxygen, have some of the highest electronegativities in the periodic table ($$\mathbf{EN = 2.5,\ 3.0,\ 3.5}$$, respectively).

The metals on the left tend to be less electronegative, with cesium having the lowest $$\mathbf{(EN = 0.7)}$$. Noble gases are excluded from this figure because they do not share electrons with other atoms since they have a full valence shell.

Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself. The larger the electronegativity value, the greater the attraction.

#### Electronegativity and Bond Type

Two idealized extremes of chemical bonding:

1. Ionic bonding-in which electrons are completely transferred from one atom to another, resulting in ions held together by electrostatic forces
2. Covalent bonding, in which electrons are shared equally between two atoms. However, most compounds have polar covalent bonds, which means that electrons are not shared equally between the bonded atoms.

A polar covalent bond’s electronegativity determines how electrons are shared between the two atoms. An atom’s electronegativity increases with its ability to attract electrons in its bonds. The electrons in a polar covalent bond are shifted toward the more electronegative atom; therefore, the more electronegative atom is the one with a partial negative charge.

The greater the difference in electronegativity, the more polarized the electron distribution and the higher the partial charges. Recall that a lowercase Greek delta $$\mathbf{δ}$$ is used to indicate that a bonded atom possesses a partial positive charge, indicated by $$\mathbf{δ^+}$$, or a partial negative charge, indicated by $$\mathbf{δ^−}$$, and a bond between two atoms that possess partial charges is a polar bond.

• Nonpolar covalent bond: Electrons are shared equally between two atoms. No charges on atoms.
• Polar covalent bond: Covalent electrons are shared unequally between atoms. Atoms have partial charges.
• Ionic bond: Complete transfer of one or more valence electrons. The resulting ions have full charges.

The Electron Distribution in a Nonpolar Covalent Bond, a Polar Covalent Bond, and an Ionic Bond Using Lewis Electron Structures. Electron-rich (negatively charged) regions are shown in blue; electron-poor (positively charged) regions are shown in red.

A bond can be classified as ionic, nonpolar covalent, or polar covalent by calculating the difference in electronegativity $$\mathbf{(Delta{EN})}$$ of two bonded atoms. If 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 $$\mathbf{H–H,\ H–Cl,\ and\ Na–Cl\ are\ 0\ (nonpolar),\ 0.9\ (polar covalent),\ and\ 2.1\ (ionic)}$$, respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). The figure below shows the relationship between electronegativity difference and bond type.

Despite the fact that this is just a general guide, there are many exceptions. When determining the covalent or ionic nature of a bond, it is best to consider the types of atoms involved and their relative positions in the periodic table. The bond between two nonmetals is usually covalent; whereas the bond between a metal and a nonmetal is usually ionic.

Pure covalent bonds have zero electronegativity between the bonding atoms. A polar covalent bond has an electronegativity difference between $$\mathbf{0.4\ and\ 1.8}$$. Ionic bonds have an electronegativity difference exceeding $$\mathbf{1.8}$$. The figure above shows As the electronegativity difference increases between two atoms, the bond becomes more ionic. Compounds may contain both covalent and ionic bonds.

Polyatomic ions, such as $$\mathbf{OH^–,\ NO_{3^−},\ and\ NH_{4^+}}$$, are held together by polar covalent bonds. By combining ions with opposite charges, these polyatomic ions can form ionic compounds. As an example, potassium nitrate, $$\mathbf{KNO3}$$, contains the $$\mathbf{K^+}$$ cation and the polyatomic $$\mathbf{NO_{3^−}}$$ anion.

In potassium nitrate, the bonds result from electrostatic attraction between the ions $$\mathbf{K^+}$$ and $$\mathbf{NO_{3^−}}$$, as well as covalent bonding between nitrogen and oxygen atoms in $$\mathbf{NO_{3^−}}$$.

#### Electronegativity and Bond Polarity

In spite of our definition of covalent bonding as electron sharing, electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, one atom will always attract electrons in the bond more strongly than the other atom, as shown in the Figure below.

When such an imbalance occurs, there is a resulting buildup of some negative charge on one side of the bond and some positive charge on the other side of the bond.

As shown in part (b) of the Figure below, a polar covalent bond has an unequal sharing of electrons. The covalent bond that has an equal sharing of electrons (part (a) of the Figure below) is a nonpolar covalent bond.

Covalent bonds between atoms of different elements are polar bonds, but their degree of polarity varies widely. A few of the bonds between different elements are minimally polar, while others are strongly polar. Electrons are transferred rather than shared in ionic bonds, which can be considered the ultimate in polarity.

Chemists use electronegativity, a measure of how strongly an atom attracts electrons when it forms a covalent bond, to determine the degree of polarity of a covalent bond. Electronegativity can be rated on a variety of numerical scales. Among the most popular is the Pauling scale.

The polarity of a covalent bond can be determined by comparing the electronegativities of the two atoms that form it. The greater the electronegativities, the greater the imbalance of electron sharing in a bond.

There are no hard and fast rules, but if the difference between electronegativities is less than about $$\mathbf{0.4}$$, the bond is considered nonpolar; if the difference is greater than $$\mathbf{0.4}$$, the bond is considered polar. The resulting compound is considered ionic rather than covalent if the difference in electronegativities is large enough $$\mathbf{(generally\ greater\ than\ 1.8)}$$. Zero electronegativity difference, of course, indicates a nonpolar covalent bond.

As a result of a molecule’s polar bonds, the molecule can display an uneven distribution of charge, depending on the orientation of its individual bonds. The orientation of the two $$\mathbf{O–H}$$ bonds in a water molecule is bent: one end of the molecule has a partial positive charge, and the other end has a partial negative charge.

The molecule itself is polar. The physical and chemical properties of water are greatly affected by its polarity. (For example, the boiling point of water $$\mathbf{[100°C]\} is high for such a small molecule, as polar molecules attract each other strongly.) On the other hand, while the two \(\mathbf{C=O}$$ bonds in carbon dioxide are polar, they are directly opposite one another and cancel each other out. This makes carbon dioxide molecules nonpolar in general. Some of carbon dioxide’s properties are affected by its lack of polarity. (For example, carbon dioxide becomes a gas at $$\mathbf{77°C}$$, almost $$\mathbf{200°C}$$ lower than the boiling point of water.)

### Properties of Polar Covalent Compounds

• Physical state: Due to their strong interactions, these compounds can exist as solids.
• Melting and boiling points: Polar compounds have higher melting and boiling points than non-polar compounds.
• Conductivity: Because the ions are mobile in the solution state, they conduct electricity.
• Solubility: These are highly soluble in polar solvents such as water.

### Dipole Moment

The polarity of a covalent bond is explained by a physical quantity called Dipole moment $$\mathbf{(μ)}$$. Dipole moments are defined as the product of charge and separation of charges. The dipole moment is denoted by $$\mathbf{‘μ’}$$ and it’s units is $$\mathbf{Debye\ (or)\ esu-cm}$$.

$$μ = (e \times d)\ esu-cm$$

• $$\mathbf{d \rightarrow distance\ between\ charge\ or\ bond\ length}$$
• $$\mathbf{e \rightarrow electronic\ charge}$$

The net bond dipole moment is

$$μ=\sqrt{{μ_1}^2+{μ_2}^2+2μ_1μ_2\cos \theta}$$

Here $$\mathbf{\theta = bond\ angle}$$

### Characteristics of Dipole Moment

• Dipole moment is a vector quantity.
• The dipole moment is zero for non-polar molecules.
• For symmetrically applicable molecular dipole moment is = 0.
• E.g. 1: Carbon dioxide being linear the net bond moment is equal to zero since the individual bond moment cancels with each other.
• E.g. 2: Carbon tetrachloride has zero dipole moment since the molecular is highly symmetrical with tetrahydro structure.
• Dipole moment is used to calculate the percentage ionic character of a covalent bond.

## Conclusion

In covalent bonds, electrons are shared between atoms and are attracted to the nuclei of both atoms. Electrons in pure covalent bonds are equally shared. The electrons in polar covalent bonds are not equally shared, as one atom exerts a stronger force of attraction on the electrons than the other. An atom’s ability to attract electrons in a chemical bond is called its electronegativity.

A bond’s polarity is determined by the difference in electronegativity between two atoms. In a diatomic molecule containing two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent. The bond is characterized as ionic when the electronegativity difference is large, as it is between metals and nonmetals.

• A non-polar covalent bond occurs when there is no difference in electronegativity between two atoms.
• A small electronegativity difference leads to polar covalent bonds.
• Large electronegativity differences lead to ionic bonds.

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## FAQs

What is an example of a polar covalent bond?

Water, abbreviated as $$\mathbf{H2O}$$, is an example of a polar covalent bond. Electrons are unequally distributed, with oxygen atoms spending more time with electrons than hydrogen atoms. The oxygen atom carries a partial negative charge since electrons spend more time with it.

Which has a polar covalent bond?

Usually, polar covalent bonds are formed between two nonmetal atoms with different electronegativity levels. Consider A and B, in which their electronegativity difference is not equal to zero and they contain a covalent bond between them.

How do you know if a covalent bond is polar or nonpolar?

Even though there are no hard and fast rules, if the difference in electronegativities is less than $$\mathbf{0.4}$$, the bond is considered nonpolar; if the difference is greater than $$\mathbf{0.4}$$, the bond is considered polar.

Is $$\mathbf{N_2}$$ a polar covalent bond?

The nitrogen atom is equally electronegative as the other atoms in the molecule, so electrons are shared equally between them. As a result, $$\mathbf{N≡N}$$ is a non-polar bond. Therefore, dinitrogen does not contain polar bonds.

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