Chemical Bonding: Definition, Types, Questions, Examples

Edited By Shivani Poonia | Updated on Jul 02, 2025 07:58 PM IST

Chemical bonding is a very basic concept of chemistry. A description associated with the attractive forces that keep together the atomic constituents in a single unit is referred to either as a molecule or compound. It deals with complex issues like electrostatic forces, quantum mechanics, and physical properties of material. Indeed, proper knowledge of a chemical bond is indispensable to provide the full exemplification of the characteristics and behavior of matter at the atomic and molecular levels. Chemical bonding is quite an intriguing field that has seen one theory or model after another regarding explaining the nature of these forces of attraction.

This Story also Contains

Cause of Chemical Combination
Lewis's symbol of elements
Theories of Chemical Bonding
Formation of Covalent Bonds
Coordinate Bonds
Types of Chemical Bonds
Applications of Chemical Bonding
Some Solved Examples
Conclusion

Chemical Bonding: Definition, Types, Questions, Examples

Cause of Chemical Combination

Atoms combine based on the following reasons:

(i) Decrease in energy: All systems in the universe tend to lose potential energy and achieve more stability. It is an observed fact that a bonded state is more stable than an unbonded state because the bonded state has lower potential energy than the unbonded state. Thus when two atoms approach each other, they combine only under the condition that there is a decrease in potential energy. When two atoms approach each other, new kinds of forces of attraction and repulsion start acting. These forces are:

Electrons and nuclei attract each other: Attractive forces are always energetically favorable, thus an electron attracted to a nucleus is of lower energy and therefore more stable than a free electron.
Electrons repel each other: Because of this repulsion, the energy is raised and the stability reduced.
Nuclei repel each other: The repulsion exists between the nuclei and this also reduces the stability.

Among all these above forces, If the net result is the attraction, then the total potential energy of the system decreases and a chemical bond formation takes place. No chemical bonding is possible if the net result is repulsion.

(ii) Lewis Octet Rule: The atoms of all elements during the bond formation try to attain the stable noble gas configuration, i.e., they try to obtain either 2 electrons (when only one energy shell) or 8 electrons in their outermost energy level which is of maximum stability and hence of minimum energy. Thus, the tendency of atoms to achieve eight electrons in their outermost shell is known as the Lewis octet rule. The octet rule is the basis of the electronic theory of valency. All the noble gases like helium, neon, etc. are not active towards the bond formation because of their already filled outermost shell, in other words, their octet is already complete and thus these elements do not need to combine with other elements to complete its octet.

Lewis's symbol of elements

For explaining the formation of bonds, the Lewis symbol representation of atoms is necessary. To write the Lewis symbol for an element, we write down its symbol surrounded by several dots that are equal to the number of valence electrons. Paired and unpaired valence electrons are also indicated. The Lewis symbols for some of the elements like Chlorine, Aluminium, and Argon are mentioned below:

Theories of Chemical Bonding

Several theories have been postulated to satisfactorily account for and describe the nature of a chemical bond formed. Each theory, however, is subject to several strengths and limitations. Some of the widely accepted theories are:
1. Lewis Theory: This theory was put forward by Gilbert N. Lewis in the year 1916. The theory underlines the valence electronic structure and the concept of sharing or transferring electrons to achieve a noble gas configuration.
2. Valence Bond Theory: This concept was derived by Linus Pauling, who focused on the interaction of atomic orbitals and the formation of localized electron pair bonds. A few of the concepts like orbital hybridization and resonance are introduced in the theory to explain the stability as well as the directional nature of the bond.

3. Molecular Orbital Theory: The work of Robert Mulliken laid down in the molecular orbital theory states that delocalized molecular orbitals are brought up by the linear combination of atomic orbitals. It explains the energetics and pattern of bonds in molecules.Such theories explain the strength, directionality, and polarity of chemical bonds. The aforesaid ingredients are of immense importance in understanding the structure and properties of matter.

Formation of Covalent Bonds

Nonmetal atoms frequently form covalent bonds with other nonmetal atoms. For example - the hydrogen molecule, H₂, contains a covalent bond between its two hydrogen atoms. The figure given below shows the explanation of this bond. Starting on the far right, we have two separate hydrogen atoms with a particular potential energy, indicated by the red line. Along the x-axis is the distance between the two atoms. As the two atoms approach each other their valence orbitals (1s) begin to overlap. The single electrons on each hydrogen atom then interact with both atomic nuclei, occupying the space around both atoms. The strong attraction of each shared electron to both nuclei stabilizes the system, and the potential energy decreases as the bond distance decreases. If the atoms continue to approach each other, the positive charges in the two nuclei begin to repel each other, and the potential energy increases. The bond length is determined by the distance at which the lowest potential energy is achieved.

Fomation of Covalent bond
The potential energy of two separate hydrogen atoms (right) decreases as they approach each other, and the single electrons on each atom are shared to form a covalent bond. The bond length is the internuclear distance at which the lowest potential energy is achieved.

Coordinate Bonds

It is a special type of covalent bond in which both the shared electrons are contributed by one atom only. Such a bond is also known as a dative bond. A coordinate or a dative bond is established between two such types of atoms, out of which one has a complete octet and while the other is short of a pair of electrons. This bond is represented by “→”.

The atom that donates the electron pair is called the donor while the atom which accepts the electron pair is called the acceptor. The compounds in which the coordinate bond exists are known as complex or coordination compounds. Some example include [Pt(en)₂]CO₃, [Ni(H₂O)₆]Cl₂, etc.

Characteristics of Coordination Compounds

The main properties of the coordination compounds are mentioned below:

Melting and boiling points: The melting and boiling points of these compounds are higher than purely covalent compounds but lower than purely ionic compounds.
Solubility: These compounds are sparingly soluble in polar solvents like water but readily soluble in non-polar solvents.
Stability: The stability of these compounds is similar to the covalent compounds.
Conductivity: Like covalent compounds, these are also bad conductors of electricity.
Dielectric constant: The compounds containing coordinate bonds have high values of the dielectric constants.

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Types of Chemical Bonds

There are many different types of chemical bonds, though they are best sorted according to what is actually occurring in the attractive forces between the atoms:

1. Ionic Bonds: Formed by a metal transferring an electron to a nonmetal, thus creating two electrically opposite ions held together by electrostatic forces.

2. Covalent Bonds: Here, both atoms combine to share electrons; this way, through electron-sharing, each of the atoms involved gets the electronic configuration of having eight electrons in its valence shell. The sharing can be equal.

3. Metallic Bonds: It is the kind of attraction that occurs between the atoms in those materials that are good conductors of electricity, such as the pure form of metals:

4. Hydrogen Bonds: Not a primary bond, hydrogen bonding is a special kind of dipole-dipole interaction that occurs when a hydrogen atom covalently bonded to an atom of high electronegativity [(nitrogen, oxygen, or fluorine)] gets close to another highly electronegative atom.

Applications of Chemical Bonding

Chemical bonding is a universal concept that underlines many day-to-day world applications. A comprehension of chemical bonding principles applies to the following:

1. Materials Science:, the physical properties that substances exhibit—such as electrical conductivity, malleability, tensile strength, and melting point—have their roots in the nature and strength of their chemical bonds. This is necessary knowledge to design and develop new materials with special characteristics.

2. Biochemistry: The stability and specificity of biomolecules, including proteins or nucleic acids, derive from the highly intricate patterns of chemical bonds created between the bonded atoms that form the biomolecule. This accounts for the reasons that such bonds have to be thoroughly understood so that one may progressively get to the root of the structural and functional aspects of life.

3. Nanotechnology: In such applications, concepts of chemical bonds are put to use in the manipulation and control of atoms and molecules. Having in mind the vision of desired functionalities, nanostructures, and nanodevices are sketched out, and small prototypes of the same are built, with selective formations and breaking of bonds.

4. Environmental Chemistry: Chemical bonding forms the critical underpinning to understanding and solving environmental problems. That is the formation of pollutants, contaminant behavior of soil and water, and material and process designs for sustainability.

5. Forensic Science: Chemical bond information can be used for forensic science purposes in the sense that it can identify unknown substances to determine the source of materials and serve in the courts as evidence for the presence of particular compounds on a crime scene.

Some Solved Examples

Example 1

Question: The molecules are made of two or more atoms joined together by some force acting between them. This force is termed a:
1) Covalent bond
2) Co-ordinate bond
3) (correct) Chemical bond
4) Ionic bond

Solution: The molecules are made of two or more atoms joined together by some force acting between them. This force is termed a chemical bond[1]. Chemical Bonds are very stable compared to other bonds. Hence, the answer is option (3).

Example 2

Question: When two atoms approach each other, they combine only under condition that there should be:
1) Increase in potential energy
2) Increase in kinetic energy
3) (correct) Decrease in potential energy
4) Increase in kinetic energy

Solution: It is a fundamental truth that all-natural systems tend to lose potential energy and become more stable. It is an observed fact that a bonded state is more stable than an unbonded state. This is because the bonded state has lower potential energy than the unbonded state. Hence, when two atoms approach each other, they combine only under the condition that there is a decrease in potential energy. Therefore, the answer is option (3).

Example 3

Question: Octet rule is based upon:
1) Shape of the molecules
2) (correct) Chemical inertness of noble gases
3) Energy of the molecule
4) None

Solution: As we have learned, the Octet rule is based on the chemical inertness of noble gases. Since the number of outermost electrons influences the chemical behavior of an atom, the number of eight so-called valence electrons (or two in the case of helium) means a particularly stable electron occupation. Based on the noble gases with eight valence electrons (except helium), the effort to achieve the noble gas configuration is also called the octet rule. Hence, the answer is option (2).

Example 4

Question: Which of the following statements is not true regarding the electronic theory of chemical bonding?
1) The theory explains the formation of a chemical bond by the sharing of electrons between two atoms
2) The theory is also known as the covalent bond theory
3) The theory assumes that atoms try to achieve stability by completing their octet or duplet
4) (correct) The theory cannot explain the formation of ionic bonds

Solution: The electronic theory of chemical bonding is based on the sharing of electrons between two atoms. It explains how the atoms in a molecule are held together by a chemical bond. This theory is also known as the covalent bond theory. According to this theory, atoms achieve stability by completing their octet or duplet configuration. However, this theory cannot explain the formation of ionic bonds, which involves the transfer of electrons from one atom to another. Hence, the answer is option (4).

Example 5

Question: According to the electronic theory of chemical bonding, which of the following is not a correct statement?
1) Atoms bond together to attain a more stable electron configuration
2) A bond is formed when two atoms have overlapping orbitals
3) (correct) A bond is formed when the electronegativity difference between two atoms is less than 1.7
4) A bond is formed when the potential energy of the system is lowered

Solution: According to the electronic theory of chemical bonding, atoms bond together to attain a more stable electron configuration. A bond is formed when two atoms have overlapping orbitals and when the potential energy of the system is lowered. However, the theory does not specify an electronegativity difference of less than 1.7 as a requirement for bond formation. Therefore, option (3) is not a correct statement according to the electronic theory of chemical bonding.

Conclusion

The creatures of the most insistent and fascinating area of chemistry still lie within the area of chemical bonding. The theories and types of chemical bonds explain the structure and properties of matter from the simplest molecules to the largest macroscopic structures. Their applications find wide areas that cut across materials science to biochemistry, hosting their key functions on our way of life and how they are upgrading the intellectual contributions made toward science.

Frequently Asked Questions (FAQs)

1. How do you determine if a molecule is polar or nonpolar?

To determine if a molecule is polar or nonpolar, consider:

2. How does bond length affect bond strength?

Generally, shorter bond lengths result in stronger bonds. This is because the closer the atoms are, the greater the overlap of their electron orbitals, leading to stronger attractive forces between the nuclei and the shared electrons.

3. How do intermolecular forces differ from chemical bonds?

Intermolecular forces are attractions between molecules, while chemical bonds are attractions within molecules. Intermolecular forces are generally weaker than chemical bonds and are responsible for properties like boiling point and solubility.

4. How do bond order and bond strength relate?

Bond order is the number of electron pairs shared between two atoms in a molecule. Higher bond orders generally result in stronger and shorter bonds. For example, a triple bond (bond order of 3) is stronger than a double bond (bond order of 2).

5. How does hybridization affect molecular geometry?

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals. It affects molecular geometry by determining the arrangement of electron pairs around the central atom, which in turn influences the shape and bond angles of the molecule.

6. Why do atoms form chemical bonds?

Atoms form chemical bonds to achieve a more stable electron configuration, typically by filling their outermost electron shell. This process lowers the overall energy of the system and increases stability.

7. What is the relationship between atomic size and bond strength?

Generally, as atomic size increases down a group in the periodic table, bond strength decreases. This is because larger atoms have their valence electrons farther from the nucleus, resulting in weaker attraction to the bonding electrons and thus weaker bonds.

8. How does electronegativity affect the distribution of electron density in a bond?

In a bond between atoms with different electronegativities, the electron density is shifted towards the more electronegative atom. This creates a partial negative charge on the more electronegative atom and a partial positive charge on the less electronegative atom, resulting in a polar bond.

9. How does the concept of orbital overlap relate to bond strength?

Greater orbital overlap generally leads to stronger bonds. The extent of overlap depends on the size and shape of the orbitals involved. For example, the strong overlap in sigma bonds contributes to their strength, while the side-by-side overlap in pi bonds results in weaker bonds.

10. How does the concept of electronegativity explain the formation of hydrogen bonds?

Hydrogen bonds form between a hydrogen atom bonded to a highly electronegative atom (like N, O, or F) and another highly electronegative atom. The high electronegativity difference creates a strong dipole in the X-H bond, allowing the partially positive H to interact with the lone pair on another electronegative atom.

11. What is a coordinate covalent bond?

A coordinate covalent bond (also called a dative bond) is a type of covalent bond where both shared electrons come from the same atom. The atom donating the electron pair is called the donor, and the atom accepting it is the acceptor.

12. How does electronegativity difference affect bond character?

The electronegativity difference between bonded atoms determines the bond character:

13. How do hydrogen bonds influence the properties of water?

Hydrogen bonds between water molecules are responsible for many of water's unique properties, including:

14. How do formal charges help in determining the most stable Lewis structure?

Formal charges help identify the most stable Lewis structure by showing the distribution of electrons in a molecule. The most stable structure typically has:

15. What is the relationship between bond length and bond energy?

Bond length and bond energy are inversely related. As bond length decreases, bond energy increases. This is because shorter bonds allow for greater overlap of atomic orbitals, resulting in stronger bonds that require more energy to break.

16. What are valence electrons, and how do they relate to chemical bonding?

Valence electrons are the electrons in the outermost shell of an atom. They are the electrons involved in chemical bonding. The number of valence electrons an atom has determines its bonding behavior and the types of bonds it can form.

17. What is resonance in chemical bonding?

Resonance is a concept used to describe delocalized electrons in molecules that cannot be accurately represented by a single Lewis structure. It involves the use of multiple structures to represent the actual electron distribution in a molecule.

18. What is a polar covalent bond?

A polar covalent bond is a type of covalent bond where electrons are shared unequally between atoms due to differences in electronegativity. This results in a partial positive charge on one atom and a partial negative charge on the other.

19. What is the difference between sigma and pi bonds?

Sigma (σ) bonds are formed by the head-on overlap of atomic orbitals along the internuclear axis. Pi (π) bonds are formed by the side-by-side overlap of p orbitals perpendicular to the internuclear axis. Sigma bonds are stronger and more flexible than pi bonds.

20. What is the role of electron-domain geometry in determining molecular shape?

Electron-domain geometry considers the arrangement of all electron pairs (bonding and non-bonding) around a central atom. This arrangement determines the molecular shape by influencing the positions of the bonded atoms while accounting for lone pair repulsions.

21. What is the role of lone pairs in determining molecular geometry?

Lone pairs occupy space around the central atom but don't contribute to the overall shape of the molecule. However, they exert a strong repulsive force on bonding pairs, often causing deviations from ideal geometries. For example, the bent shape of water is due to the presence of two lone pairs on the oxygen atom.

22. How does the octet rule explain the formation of polyatomic ions?

The octet rule drives the formation of polyatomic ions as atoms combine to achieve stable electron configurations. For example, in the carbonate ion (CO₃²⁻), the central carbon shares electrons with three oxygen atoms and gains a pair from one oxygen, allowing all atoms to achieve full octets while creating a -2 charge.

23. How does the concept of hybridization explain the tetrahedral structure of methane?

In methane (CH₄), the carbon atom undergoes sp³ hybridization. This creates four equivalent hybrid orbitals arranged tetrahedrally around the carbon atom. These orbitals overlap with the 1s orbitals of the four hydrogen atoms, resulting in the tetrahedral structure of methane.

24. What is the relationship between bond order and bond length?

As bond order increases, bond length generally decreases. This is because higher bond orders involve more shared electrons, leading to stronger attraction between the nuclei and a shorter distance between them. For example, C-C < C=C < C≡C in terms of bond length.

25. How does the concept of resonance explain the strength of the carbon-oxygen double bond in carbon dioxide?

The strength of the C=O bonds in CO₂ is explained by resonance. CO₂ has two equivalent resonance structures, each with one double and one single C-O bond. The actual structure is a hybrid of these, resulting in two equivalent C-O bonds with a bond order of 2, stronger than a typical single bond but weaker than a typical double bond.

26. What is the octet rule, and why is it important in chemical bonding?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons (like noble gases). This rule helps predict how atoms will bond and form molecules, although there are exceptions.

27. How does the VSEPR theory help predict molecular geometry?

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry based on the arrangement of electron pairs around a central atom. It states that electron pairs repel each other and will adopt positions that minimize these repulsions, determining the molecule's shape.

28. How does electron affinity affect bond formation?

Electron affinity is the energy change when an atom gains an electron. Elements with high electron affinities (like halogens) are more likely to form ionic bonds by accepting electrons from elements with low ionization energies (like alkali metals).

29. What are antibonding orbitals, and how do they affect molecular stability?

Antibonding orbitals are high-energy molecular orbitals that, when occupied, weaken the bond between atoms. The presence of electrons in antibonding orbitals decreases molecular stability and can lead to bond breaking.

30. What is the difference between bond energy and bond enthalpy?

Bond energy is the energy required to break a specific bond in a molecule, while bond enthalpy is the average energy required to break all bonds of a particular type in various molecules. Bond enthalpy is more commonly used as it's easier to measure experimentally.

31. What is chemical bonding?

Chemical bonding is the process by which atoms join together to form molecules or compounds. It involves the sharing or transfer of electrons between atoms to achieve a more stable electron configuration.

32. What are the three main types of chemical bonds?

The three main types of chemical bonds are:

33. How does electronegativity affect bond formation?

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Differences in electronegativity between atoms determine the type of bond formed. Large differences lead to ionic bonds, while small differences result in covalent bonds.

34. What is an ionic bond?

An ionic bond is a type of chemical bond formed by the complete transfer of electrons from one atom (usually a metal) to another atom (usually a non-metal). This results in the formation of oppositely charged ions that are held together by electrostatic attraction.

35. How does a covalent bond differ from an ionic bond?

In a covalent bond, electrons are shared between atoms, while in an ionic bond, electrons are transferred from one atom to another. Covalent bonds typically form between non-metals, while ionic bonds form between metals and non-metals.

36. What is the relationship between bond polarity and molecular polarity?

Bond polarity contributes to molecular polarity, but molecular geometry also plays a crucial role. A molecule with polar bonds can be nonpolar if the bond dipoles cancel each other out due to the molecule's symmetrical geometry.

37. How do metallic bonds differ from ionic and covalent bonds?

Metallic bonds involve the sharing of delocalized electrons among a lattice of positive metal ions. Unlike ionic or covalent bonds, which involve specific atoms, metallic bonds create a "sea" of electrons that can move freely throughout the metal structure.

38. What is the relationship between bond angle and hybridization?

The type of hybridization determines the bond angles in a molecule:

39. What is the difference between a single, double, and triple covalent bond?

The difference lies in the number of electron pairs shared between atoms:

40. What is the difference between sigma and pi bonds in terms of rotation?

Sigma (σ) bonds allow free rotation around the bond axis because the electron density is symmetrical around this axis. Pi (π) bonds restrict rotation because the electron density is above and below the bond axis, and rotation would require breaking and reforming the bond.

41. How does the concept of resonance explain the structure of benzene?

Resonance in benzene explains its unique stability and structure. The true structure of benzene is a hybrid of two resonance forms, each with alternating single and double bonds. This results in all carbon-carbon bonds being equivalent, with a bond order of 1.5 and electron delocalization throughout the ring.

42. What is the difference between a coordinate covalent bond and a regular covalent bond?

In a coordinate covalent bond, both shared electrons come from one atom (the donor), while in a regular covalent bond, each atom contributes one electron. However, once formed, coordinate covalent bonds behave similarly to regular covalent bonds and cannot be distinguished in the final molecule.

43. How does bond polarity contribute to the solubility of substances?

Bond polarity affects molecular polarity, which in turn influences solubility. Polar molecules tend to dissolve in polar solvents (like water) due to favorable dipole-dipole interactions. Nonpolar molecules, with nonpolar bonds, tend to be insoluble in polar solvents but soluble in nonpolar solvents.

44. What is the difference between bonding and antibonding molecular orbitals?

Bonding molecular orbitals have lower energy than the original atomic orbitals and increase electron density between nuclei, strengthening the bond. Antibonding molecular orbitals have higher energy than the original atomic orbitals and decrease electron density between nuclei, weakening the bond.

45. How does the concept of hybridization explain the planar structure of ethene (C₂H₄)?

In ethene, each carbon atom undergoes sp² hybridization, creating three sp² hybrid orbitals in a trigonal planar arrangement. These orbitals form sigma bonds with two hydrogens and the other carbon. The remaining unhybridized p orbital on each carbon forms a pi bond, resulting in the planar structure of ethene.

46. What is the relationship between bond polarity and dipole moment?

Bond polarity contributes to a molecule's dipole moment, but the overall molecular geometry also plays a crucial role. A molecule with polar bonds may have a zero dipole moment if the bond dipoles cancel each other out due to symmetry (e.g., CO₂). The net dipole moment is the vector sum of all bond dipoles in the molecule.

47. What is the relationship between bond energy and reaction enthalpy?

The enthalpy change of a reaction can be estimated by comparing the total bond energies of the reactants and products. Breaking bonds requires energy (endothermic), while forming bonds releases energy (exothermic). The overall enthalpy change is the difference between the energy required to break reactant bonds and the energy released in forming product bonds.

48. What is the relationship between bond angle and molecular polarity?

Bond angle can significantly affect molecular polarity. In molecules with polar bonds, a bent geometry (like H₂O) results in a net dipole moment, while a linear geometry (like CO₂) can result in dipole cancellation and a nonpolar molecule. The bond angle determines how the individual bond dipoles combine to produce the overall molecular polarity.

49. How does the concept of hybridization explain the structure of acetylene (C₂H₂)?

In acetylene, each carbon atom undergoes sp hybridization, creating two sp hybrid orbitals in a linear arrangement. These form a sigma bond between the carbons and with the hydrogens. The remaining two unhybridized p orbitals on each carbon form two pi bonds, resulting in the linear structure and triple bond of acetylene.

50. What is the relationship between bond type and physical properties of substances?

The type of bonding in a substance greatly influences its physical properties: