Covalent Bond: When two atoms’ valence shell electrons are shared, a covalent bond is formed. It is in charge of binding a single molecule’s atoms together. When electrical things are damp, we’re trained not to contact them with our naked hands. Elements with extremely high ionisation energies cannot transport electrons, while those with extremely low electron affinity cannot absorb them. 

The atoms of such elements tend to share electrons with atoms of other elements or atoms of the same element, resulting in both atoms achieving an octet configuration in their respective valence shells and therefore achieving stability. A “covalent bond” is a relationship produced by the exchanging of electron pairs between molecules of different or similar types. Read on to know more about the covalent bond definition, covalent bond examples, and covalent compounds examples.

Covalent Bond Definition

Covalent bonding, in simple words, is the sharing of electrons between atoms to attain the noble gas configuration of the participating individual atoms.  The atoms in a covalent bond are held together by the electrostatic force of attraction. This force is in between the positively charged nuclei of the bonded atoms and the negatively charged electrons they share. 

The electrons that join atoms in a covalent bond are called the bonding pair of electrons. These bonding pair of electrons results in the formation of a discrete group of atoms called a molecule—the smallest part of a compound that retains the chemical identity of that compound. This type of bonding occurs between two atoms of the same element or of elements close to each other in the periodic table. This bonding occurs primarily between non-metals; however, it can also be observed between non-metals and metals.

What are 3 Types of Covalent Bond: Types of Chemical Bonding

The formation of covalent bonds results in a particular bond length. This bond length is a characteristic property of compounds that represents a balance between several forces such as: 

(1) the attraction between positively charged nuclei and negatively charged electrons,
(2) the repulsion between two negatively charged electrons, and 
(3) the repulsion between two positively charged nuclei.

The shorter bond length has greater bond strength. Based on the bond length, covalent bonds are of the following types.

Single Covalent Bonds Between Different Atoms

The simplest covalent bond exists in the diatomic hydrogen molecule. Halogens also exist as diatomic gases by forming covalent bonds, such as chlorine. Nitrogen and oxygen also exhibit covalent bonding by forming diatomic molecules.

Covalent bonding in molecular substances is represented by the Lewis electron dot diagram. For example, the Lewis diagram of two hydrogen atoms is:

When two hydrogen atoms share electrons, the lewis diagram is as shown below:

This depiction of molecules sharing electrons is shown by using a dash. This dash represents a covalent bond. The hydrogen molecule can be represented as:

Another element whose atoms share electrons to bond together to form diatomic (two-atom) molecules is fluorine. The Lewis diagrams of two fluorine atoms are:

Each fluorine atom completes its valence shell by contributing one valence electron. This leads to the formation of a single bond and giving each atom a complete valence shell.

The image above shows that each fluorine atom completes its valence shell and fulfils the octet rule. Like hydrogen, the fluorine molecule is also represented with a dash as shown below-

Each fluorine atom has three pairs of electrons or six electrons, which do not participate in the covalent bonding. These electrons are not shared; hence they are considered to belong to a single atom. As they are not available for bonding with other atoms, these are called non-bonding pairs (or lone pairs) of electrons.

The bond formation in hydrogen and fluorine molecules takes place between the same atoms. Let us take the examples of molecules having covalent bonds between atoms of different elements.

For example, a molecule of hydrogen fluoride comprising of a hydrogen atom and a fluorine atom:

Each atom of hydrogen and fluorine needs one additional electron to complete its valence shell. Sharing each of these electrons results in the formation of the following molecule:

In the above molecule, there are no nonbonding electrons in the hydrogen atom, while the fluorine atom has six or three pairs of nonbonding electrons (three lone electron pairs). 

Similarly, larger molecules are constructed in a similar fashion with the formation of more than one covalent bond.

Covalent Bonding in Water

The water molecule consists of an oxygen atom and two hydrogen atoms. Both of them are non-metals. Oxygen present in group \(6\) of the periodic table has  \(6\) electrons in its outer shell. Hydrogen is present in group \(1\) of the periodic table atom \(1\) electron in its outer shell. Hydrogen can share only one electron and form a single bond only.

Each of the two hydrogen atoms shares \(1\) electron present in their valence shell with oxygen to form two covalent bonds and make a water molecule \(\left( {{{\rm{H}}_2}{\rm{O}}} \right)\) as shown below-

There are \(8\) electrons present in the outer shell of the oxygen atom following the octet rule, whereas the hydrogen atom can count \(2\) electrons in its outer shell following the duplet rule. The outer shells of both oxygen and hydrogen in the water molecule have electrons close to their corresponding inert gases. This makes \({{{\rm{H}}_2}{\rm{O}}}\) molecules stable, and they will not react further with other oxygen or hydrogen atoms.

There are two pairs or \(4\) shared pairs of electrons between the atoms. Each electron pair represents one single bond-forming a single covalent bond. Two single covalent bonds are present in the water molecule.

The structural formula of a water molecule is written

Multiple Covalent Bonds

The octet rule will not be satisfied in many molecules if each pair of bonded atoms shares only two electrons. Multiple pairs of electrons need to be shared in order to obtain a stable electronic configuration.

Double Covalent Bonds Between Different Atoms

For example, carbon dioxide \(\left( {{\rm{C}}{{\rm{O}}_2}} \right).\) Each oxygen atom in a carbon dioxide molecule shares one electron with the carbon atom, as shown by the Lewis structure given below:

This completes the octet configuration neither for oxygen nor for carbon; The oxygen atom only has seven electrons in its valence shell, whereas the carbon atom only has six electrons in its valence shell. This results in the formation of an unstable molecular conformation.

In carbon dioxide, each oxygen atom shares a second electron with the central carbon atom, and the carbon atom shares one more electron with each oxygen atom:

In this arrangement, the carbon atom shares four electrons (two pairs) with the oxygen atom on the left and four electrons with the oxygen atom on the right. There are now eight electrons around each atom. Two pairs of electrons shared between two atoms make a double bond between the atoms, which is represented by a double dash:

Triple Covalent Bonds Between Different Atoms

Some molecules in which three pairs of electrons are shared by two atoms contain triple bonds. A simple compound that has a triple bond is nitrogen. Nitrogen is non-metal. Nitrogen belongs to group \({\rm{V}}\left( {\rm{A}} \right)\) of the periodic table and has \(5\) electrons in its outermost shell.

In order to attain the noble gas configuration, one nitrogen atom will share three electrons present in its valence shell with another nitrogen atom. This results in the formation of a nitrogen molecule \(\left( {{{\rm{N}}_2}} \right)\) with three covalent bonds.

By sharing six electrons (\(3\) pairs ), each nitrogen atom can count \(8\) electrons in its outer shell. These full outer shells with their shared electrons are now stable, and the \({{{\rm{N}}_2}}\) molecule will not react further with other nitrogen atoms.

Each electron pair represents one single bond. Hence, \(3\) electron pairs form \(3\) bonds called triple bonds. The triple bond is very strong, and this is what makes nitrogen so unreactive (stable).

The structural formula of a nitrogen molecule is written as follows-

\({{{\rm{N}}_2}}\) is a neutral molecule, i.e. there is no ion present (no + or – charges) in the nitrogen gas. This is because the electrons are shared, not transferred from one atom to another.

Although covalent bonding involves electron sharing, the two bonded atoms do not share the electrons equally. There will always be one atom that attracts the electrons in the bond more strongly than the other atom does unless the bond connects two atoms of the same element.

Consequently, it builds some negative charge (called a partial negative charge and (designated \({{\rm{\delta }} – }\)) on one side of the bond and some positive charge (designated as \({{\rm{\delta }} + }\)) on the other side of the bond. The ability of an atom to attract electrons in the presence of another atom is a measurable property called electronegativity and will produce a dipole moment.

How to judge the polarity of a molecule?

Through electronegativity, we can judge the relative polarity of a covalent bond. Electronegativity is a relative measure of how strongly an atom attracts electrons when it forms a covalent bond.

The most popular scale to measure electronegativity is the Pauling scale. On the Pauling scale, fluorine has the highest value, i.e. \(4\).

In a molecule, the bond type and its polarity can be predicted by taking the difference between the electronegativity values for each of the atoms involved in the bond. The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond.

A covalent bond that has an unequal sharing of electrons and the electronegativity difference is within the range \(0.1-2\) is called a polar covalent bond. 

A covalent bond that has an equal sharing of electrons and the electronegativity difference is zero is called a nonpolar covalent bond.

Polar Covalent Bond

When the electrons spend more time around the more non-metallic atom, the sharing of the electron pair becomes unequal and results in the formation of polar covalent bonds. In such a bond, there is a charge separation, with one atom being slightly more positive and the other being more negative. When atoms with different electronegativities share electrons, a polar covalent bond is formed.

For example – In the hydrogen chloride \(\left( {{\rm{HCl}}} \right)\) molecule, the atom of hydrogen and chlorine requires one more electron to form an inert gas electronic configuration. Though chlorine has a higher electronegativity than hydrogen, its atom’s attraction for electrons is not sufficient to remove an electron from the hydrogen atom.

Consequently, the bonding electrons in hydrogen chloride are shared unequally, resulting in a polar covalent bond between hydrogen and chlorine atom. This unequal sharing of the bonding pair of electrons results in the formation of a partial negative charge \(\left( {{\rm{\delta – }}} \right)\) on the chlorine atom and a partial positive charge \(\left( {{\rm{\delta + }}} \right)\)on the hydrogen atom. The symbol \({\rm{\delta }}\) (Greek lowercase delta) denotes these fractional charges.

The larger the differences between the electronegativities of the bonded atoms, the higher is the polarity of the bonds present between them. The dipole is represented by an arrow with a cross at one end. The cross is near the end of the molecule that is partially positive, and the arrowhead is near the partially negative end of the molecule.

It is due to the presence of positive and negative charges in the polar covalent molecules that enables it to interact with dipoles in other molecules. This results in dipole-dipole intermolecular forces between the molecules.

Nonpolar Covalent Bond

When electrons are equally shared between the combining atoms, a nonpolar covalent bond is formed. This phenomenon happens when there is no difference in the electronegativities of the two combining atoms. That is, to say, identical pairs of atoms form a nonpolar covalent bond. The bond is nonpolar If the electronegativity difference is less than \(0.5.\) Although oxygen is very electronegative, \({{{\rm{O}}_2}}\) is not Polar. This is because both atoms have the same electronegativity, and electrons are shared equally between them.

Given below is the list of compounds having nonpolar bonds.

Compound Name	Molecular formula
Hydrogen	\({{\rm{H}}_2}\)
Chlorine	\({\rm{C}}{{\rm{l}}_2}\)
Bromine	\({\rm{B}}{{\rm{r}}_2}\)
Iodine	\({{\rm{I}}_2}\)
Oxygen	\({{\rm{O}}_2}\)
Nitrogen	\({{\rm{N}}_2}\)

Examples of Nonpolar Covalent Molecules with Polar Covalent Bonds

Some nonpolar molecules also contain polar covalent bonds. The orientation of the various polar bonds in these molecules is such that their polarities cancel each other. Below is a list of such molecules with their formulae.

Compound Name	Molecular formula	Polar covalent bond
Carbon dioxide	\({\rm{C}}{{\rm{O}}_2}\)	\({\rm{C}} = {\rm{O}}\)
Sulfur trioxide	\({\rm{S}}{{\rm{O}}_3}\)	\({\rm{S}} = {\rm{O}}\)
Silicon dioxide	\({\rm{Si}}{{\rm{O}}_2}\)	\({\rm{Si}} = {\rm{O}}\)
Methane	\({\rm{C}}{{\rm{H}}_4}\)	\({\rm{C}} – {\rm{H}}\)
Carbon tetrachloride	\({\rm{CC}}{{\rm{l}}_4}\)	\({\rm{C}} – {\rm{Cl}}\)
Benzene	\({{\rm{C}}_6}{{\rm{H}}_6}\)	\({\rm{C}} – {\rm{H}}\)

Difference Between Polar and Nonpolar Covalent Bond

Water is a small molecule. Still, it has a high boiling point \(\left[ {100^\circ \,{\rm{C}}} \right].\) This is because water is a polar molecule. On the contrary, the two \({\rm{C}} = {\rm{O}}\) bonds in carbon dioxide lie directly opposite to each other in the molecule and cancel out each other’s effects.

However, the individual \({\rm{C}} = {\rm{O}}\) bonds are polar; as a whole, carbon dioxide molecules are nonpolar. This lack of polarity influences some of carbon dioxide’s properties. (For example, carbon dioxide becomes a gas at \( – 77^\circ \,{\rm{C,}}\) almost \(220^\circ \) lower than the temperature at which water boils.)

	Polar	Nonpolar
Type Of Atoms	Between two non-metals with different electronegativities	Between two non-metals with the same electronegativities
Electronegativity difference	\(0.1-2\)	\(0\)
Electron distribution	Asymmetrical. Unequal sharing of electrons	Symmetrical. Equal sharing of electrons
Displacement Of Shared Electrons	This bond tends to attract the bonded electrons towards the more electronegative atom, making that part negative	No displacement. Electrically neutral
Dipole Moment	Non-zero	zero
Presence of other types of bonds in compounds	Hydrophilic	Hydrophobic
Affinity towards water	High	Low
Physical properties of the compounds	High melting and boiling points	High melting and boiling points
Volatility of the liquids	Low	high
Solubility of compounds	Soluble in polar solvents	Soluble in nonpolar solvents
Examples	Water, ammonia, hydrogen chloride	Hydrogen, oxygen, nitrogen

Is a Coordinate Bond a Covalent Bond?

A coordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom. 

In the formation of a simple or ordinary covalent bond, each atom supplies at least one electron to the formation of the bond – but that is not the case every time. In the case of a coordinate covalent bond, one atom supplies both of the electrons, and the other atom does not supply any of the electrons. 

The formation of a coordinate covalent bond can be easily demonstrated by the reaction between ammonia and hydrochloric acid, which is as follows.

The reaction between ammonia and hydrochloric acid

The overall reaction is-

\({\rm{N}}{{\rm{H}}_3}\left( {\rm{g}} \right)\; + \;{\rm{HCl}}\left( {\rm{g}} \right)\; \to \;{\rm{N}}{{\rm{H}}_4}{\rm{Cl}}\left( {\rm{s}} \right)\)

The electron and proton movement during the reaction is as shown below.

When the ammonium ion, \({\rm{NH}}_4^ + ,\) is formed-

1. The hydrogen atom shown in red is the fourth hydrogen atom which is attached by a coordinate covalent bond to the ammonia molecule. As hydrogen has only one electron in its outermost shell, only the hydrogen’s nucleus is transferred from the chlorine to the nitrogen. 
2. The electron of the hydrogen atom is left behind on the chlorine to form a negative chloride ion. 
3. Once the ammonium ion has been formed, no difference lies between the coordinate covalent and the ordinary covalent bonds. All of the hydrogens are equivalent in the molecule, and the extra positive charge carried by the hydrogen ion is distributed throughout the molecule.
4. Although the electrons are shown differently in the diagram, there is no difference between them in reality. In simple diagrams, a coordinate bond is shown by a curved arrow—the arrow points from the atom, donating the lone pair to the atom accepting it.

Coordinate covalent bonds are also present in neutral molecules such as \({{\rm{O}}_2}.\)

Each oxygen atom in \({{\rm{O}}_2}\) contributes \(1\) electron to each of the bonding pairs of electrons. This results in the formation of \(2\) ‘normal’ covalent bonds or a double covalent bond \({\rm{O}} = {\rm{O}}.\)

Besides the \(2\) bonding pairs of electrons, each oxygen atom also has two nonbonding pairs of electrons called the lone pairs of electrons.

In an ozone molecule, \({{\rm{O}}_3},\) an oxygen molecule, \({{\rm{O}}_2},\) reacts with an oxygen atom in the atmosphere.
Each oxygen atom in the \({{\rm{O}}_2}\)  molecule has the stable electronic configuration of the noble gas Neon. So, the only way the third oxygen atom could add to the oxygen atoms in \({{\rm{O}}_2}\) is that one of the oxygen atoms presents in \({{\rm{O}}_2}\) to contribute both the non bonded pair of electrons to the new oxygen atom. This results in the formation of a coordinate covalent bond between the existing \({{\rm{O}}_2}\) molecule and the new oxygen atom:

In the diagram given alongside, one of the oxygen atoms from the original \({{\rm{O}}_2}\)  molecule shown in black has contributed both of its non bonded pair of electrons to the blue oxygen atom.
The blue oxygen atom has not contributed any of its existing electrons to this bonding, so this is a coordinate covalent bond (dative bond).

However, once the coordinate covalent bond has been formed, it is in no way different from a ‘normal’ covalent bond.

By looking at this Lewis structure for \({{\rm{O}}_3}\), we could say that the single bond will always be the coordinate covalent bond.

Properties of Covalent Bond

1. Most covalent compounds have relatively low melting points and boiling points.
Covalent compounds generally have low boiling and melting points due to the presence of weak intermolecular forces of attraction. These compounds are found in all three physical states at room temperature. While covalent bonds between atoms are quite strong, attractions between molecules/compounds or intermolecular forces are relatively weak.

Covalent bonds create molecules that can separate from each other when a lower amount of energy is added to them. Therefore, these compounds are highly volatile.

2. Covalent compounds usually have lower enthalpies of fusion and vaporisation than ionic compounds.
The enthalpy of fusion is the amount of energy needed, at constant pressure, to melt one mole of a solid substance. The enthalpy of vaporisation is the amount of energy, at constant pressure required to vaporise one mole of a liquid.

On average, it takes only \(1%\) to \(10%\) as much heat to change the phase of a molecular covalent compound as it does for an ionic compound.

3. Covalent compounds tend to be soft and relatively flexible.
Covalent compounds contain weak intermolecular forces of attraction, which cause these compounds to take the form of gasses, liquids, and soft solids. As with many properties, there are exceptions, primarily when molecular compounds assume crystalline forms.

4. Covalent compounds tend to be more flammable than ionic compounds.
Most flammable substances contain hydrogen and carbon atoms. These compounds can readily undergo combustion reactions in the presence of oxygen to produce carbon dioxide and water. Carbon and hydrogen have comparable electronegativities, so they are found together in many molecular compounds.

5. When dissolved in water, covalent compounds do not conduct electricity. 
Ions are needed for the passage of electricity in an aqueous solution. When mixed with water, molecular compounds dissolve into molecules rather than dissociate into ions. There are no mobile ions to conduct electricity, so they typically do not conduct electricity very well when dissolved in water.

6. Many covalent compounds are insoluble in water.
Polar covalent molecules dissolve well in a polar solvent, such as water. Examples of molecular compounds that dissolve well in water are sugar and ethanol. However, nonpolar covalent molecules do not dissolve well in water—for example – water and oil. Water cannot hydrate these molecules.

7. Many covalent compounds dissolve well in organic solvents.
Both covalent molecules and organic molecules in organic solvents are both held together by weak intermolecular forces of attraction. As they have the same type of weak intermolecular forces of attraction, the covalent molecules in the covalent compounds are easily miscible with the organic molecules in the organic solvents. Hence, covalent compounds are usually soluble in organic solvents.

8. Lewis’s theory also accounts for bond length; the stronger the bond and the more electrons shared, the shorter the bond length is.

Covalent Bond Vs Ionic Bond

	Covalent Bond	Ionic Bond
What is it?	It is a form of chemical bonding between two non-metallic atoms, which is characterised by the sharing of pairs of electrons between atoms and other covalent bonds.	Also known as an electrovalent bond, it is a type of bond formed from the strong electrostatic force of attraction between oppositely charged ions in a chemical compound.
Occurs between	Two non-metals or a non-metal and a metalloid	One metal and one non-metal
Nature of component particles	It consists of electrically neutral discrete molecules	Ionic compounds are composed of oppositely charged particles called cations and anions.
State at room temperature	Gases, liquids, or low melting solids	Crystalline solids
Polarity	Low	High
Solubility	Polar covalent compounds dissolve in polar solvents. Ex- \(\left( {{\rm{HCl}}} \right)\) in water. Nonpolar covalent compounds are soluble in organic liquids only.	Being polar in nature, ionic compounds are soluble in polar solvents only like water. Not soluble in organic liquids
Formation	They are formed between two non-metal having similar electronegativities. Neither of the atoms is strong enough to attract electrons from the other. For stabilisation, they share their electrons from the outermost orbital.	They are formed between a metal (+ion) and a non-metal (-ve ion). Non-metals are stronger than metals and can get electrons very easily from metals. These two opposite ions attract each other and form an ionic bond.
Conductivity	Nonpolar covalent compounds do not dissociate into ions. Hence, do not conduct electricity. Polar covalent compounds easily dissociate in water and behave as good conductors of electricity. Ex- \(\left( {{\rm{HCl}}} \right)\) in water.	Ionic compounds conduct electricity when molten or dissolved in water.
Melting Point	Low	High
Boiling Point	Low	High
Examples	Methane, Hydrochloric acid	Sodium Chloride, Sulphuric acid

Summary

Compounds containing covalent bonds are part and parcel of our day-to-day life. From the water used to boil an egg to the protein present inside it, all are compounds having a covalent bond. From cooking gas to the sugar in lemonade, from the oxygen we inhale to the exhalation of carbon dioxide, all consist of compounds containing covalent bonds.

The elixir of life, water, also contains covalent bonding. Hence, to summarise, it is indeed needless to say that without covalent compounds, our existence is impossible.

FAQs on Covalent Bond

Q1. Comment on the type of bond formed between two atoms if the difference in electronegativities is small? Medium? Large?
Ans: If the difference in electronegativities of the combining atoms is small, then the nonpolar covalent bond is formed. If the difference in electronegativities of the combining atoms is medium, then a nonpolar covalent bond is formed. If the difference in electronegativities of the combining atoms is large, then the bond is formed ionic.

Q2.. Is Coca Cola ionic or covalent?
Ans: Coca Cola contains a chemical known as Potassium Sorbate. The cation is \({{\rm{K}}^{ + 1}}\) and the anion is \({\left( {{{\rm{C}}_6}{{\rm{H}}_7}{{\rm{O}}_2}} \right)^{ – 1}}.\) The full chemical makeup is \({{\rm{K}}{{\rm{C}}_6}{{\rm{H}}_7}{{\rm{O}}_2}.}\) This chemical is ionic.

Q3. Why \({\rm{Be}}{{\rm{H}}_2}\) has zero dipole moment while \({\rm{BeH}}\) bond is polar?
Ans: Although the \({\rm{Be – H}}\) bonds are polar, having the same dipole moment, but due to the linear structure of \({\rm{Be}}{{\rm{H}}_2},\) the bond dipoles of the two \({\rm{Be – H}}\)  bonds cancel each other. Therefore, the resultant dipole moment of  \({\rm{Be}}{{\rm{H}}_2}\)molecule is zero.

Q4. Why does a diamond have such a high melting point?
Ans: Diamond is an allotrope of carbon. In a diamond, each of its carbon atoms is covalently bonded to four other carbon atoms. This results in the formation of a giant covalent structure. Hence, diamond is very hard and has a high melting point.

Q5. Comment on the electronegativity difference between each pair of atoms and the resulting polarity (or bond type).
Ans: \({\rm{H}}\) and \({\rm{H,}}\,{\rm{b}}{\rm{.}}\,{\rm{Na}}\) and \({\rm{Cl}}\)

The electronegativity value for hydrogen atom is \(2.1.\) Hence the difference between its electronegativities is zero resulting in nonpolar covalent bond. The electronegativity difference between sodium \(\left( {0.9} \right)\) and chlorine \(\left( {3.0} \right)\) is \(2.1,\) which is rather high, and so sodium and chlorine form an ionic compound.