Covalent Bond: Learn Definition, Types, Properties (2023)

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 amolecule—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 diatomichydrogen 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 callednon-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. Bothof them are non-metals. Oxygenpresent ingroup \(6\)of theperiodic table has \(6\)electronsin itsouter shell. Hydrogen is present in group \(1\) of the periodic table atom\(1\)electronin itsouter shell. Hydrogencan share only one electron and forma single bond only.

Eachof the twohydrogen atomsshares\(1\) electron present in their valence shell withoxygento formtwocovalent bonds and make awatermolecule\(\left( {{{\rm{H}}_2}{\rm{O}}} \right)\) as shown below-

There are\(8\) electronspresent 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 notreactfurther with otheroxygenorhydrogen atoms.

There aretwo pairsor \(4\) sharedpairs of electrons between theatoms. Eachelectron pairrepresents one single bond-forming a singlecovalent bond. Two single covalent bonds are present in the water molecule.

Thestructural formulaof awater moleculeis 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. Nitrogenisnon-metal. Nitrogenbelongs to group \({\rm{V}}\left( {\rm{A}} \right)\)of theperiodic table and has \(5\) electrons in its outermost shell.

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

Bysharingsix electrons (\(3\) pairs), eachnitrogenatomcan count\(8\) electronsin itsouter shell. Thesefullouter shellswith theirshared electronsare nowstable, and the\({{{\rm{N}}_2}}\)moleculewillnotreactfurther with othernitrogen atoms.

Eachelectron pairrepresents one single bond. Hence, \(3\) electron pairs form \(3\) bonds called triple bonds. Thetriplebondisvery strong, and this is what makesnitrogensounreactive(stable).

Thestructural formulaof anitrogen moleculeis written as follows-

\({{{\rm{N}}_2}}\) is a neutral molecule, i.e. there isnoionpresent (no+or–charges) inthe nitrogen gas. This is because theelectronsareshared, nottransferredfrom oneatomtoanother.

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 calledelectronegativity 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 apolar covalent bond.

A covalent bond that has an equal sharing of electrons and the electronegativity difference is zero is called anonpolar 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 thehydrogen chloride\(\left( {{\rm{HCl}}} \right)\) molecule, the atom of hydrogen and chlorine requires one more electron to form an inert gaselectronic configuration. Though chlorine has a higherelectronegativitythan 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 apolar 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 thechlorine atomand a partial positive charge \(\left( {{\rm{\delta + }}} \right)\)on thehydrogen 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 isnonpolar 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.

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.

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 moleculeand 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.)

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 whichbothelectrons 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 inblackhas contributed both of its non bonded pair of electrons to theblueoxygen atom.
Theblueoxygen 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 bondscreate 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 andvaporisation than ionic compounds.
The enthalpy of fusionis the amount of energy needed, at constant pressure, to melt one mole of a solid substance.The enthalpy of vaporisationis 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 thephase of a molecular covalent compoundas it does for an ionic compound.

3. Covalent compoundstend to be soft and relatively flexible.
Covalent compounds contain weak intermolecular forces of attraction, which cause these compounds to takethe form of gasses, liquids,and soft solids. Aswith 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 producecarbon dioxideand 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 free 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 moleculesdissolve 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 byweak 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 solublein 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

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 Colacontains 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 isionic.

Q3. Why \({\rm{Be}}{{\rm{H}}_2}\) has zero dipole moment while \({\rm{BeH}}\) bond is polar?
Ans: Althoughthe \({\rm{Be – H}}\)bonds are polar, having the samedipole moment, but due to the linear structure of\({\rm{Be}}{{\rm{H}}_2},\) thebond dipolesof the two \({\rm{Be – H}}\) bondscancel each other. Therefore, the resultantdipole momentof \({\rm{Be}}{{\rm{H}}_2}\)moleculeiszero.

Q4. Why does a diamond have such a high melting point?
Ans: Diamond isan allotrope of carbon. In a diamond, each of its carbon atomsiscovalently bonded to four other carbon atoms. This results in the formation of a giant covalent structure. Hence,diamond isvery hard andhasahigh 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.