Without a firm grip on the concept of chemical bonding, you cannot fully comprehend chemistry as a subject. During a chemical reaction, a new molecule is formed by the attraction of atoms or ions. A strong, attractive force that holds two or more atoms together in an independent entity, i.e., a stable molecule, is known as a chemical bond.
This introductory article will discuss everything you need to know about chemical bonding, including basic terms about the atomic structure, different types of chemical bonds, their definitions and examples. You will also learn about the different theories of chemical bonding.
So, without any further delay, dive into the article and let’s start reading!
What is chemical bonding -Definition
A chemical bond is defined as a force of attraction that holds atoms and/or ions together in a stable molecule or a crystal structure. The process of developing this strong, attractive force or association between the constituents of a molecule is known as chemical bonding.
History and origin of chemical bonding
The concept of chemical bonding came into existence with Sir Issac Newton’s idea of an attractive force back in 1709. It gained prominence in the years that followed with the creation of the Affinity Table by Geoffroy (a French chemist) in 1718 and the Lewis dot diagram introduced by Gilbert N. Lewis (an American chemist) in 1916. In the modern scientific world today, we can apply vibrational spectroscopic techniques such as IR spectroscopy to affirm the presence of a definite chemical bond between constituent atoms.
Basic concepts related to atomic structure
An atom is the smallest, indivisible part of an element. It comprises a dense nucleus at the center. The nucleus is positively charged as it consists of positively charged sub-atomic particles called protons and neutral particles called neutrons. In contrast, the negatively charged sub-atomic particles (electrons) circulate the nucleus in specific orbits called energy levels or shells. A shell consists of different atomic orbitals. The atomic orbital is a region exhibiting the maximum probability of finding an electron.
A particular force of attraction exists between the oppositely charged protons and electrons that keep the atomic structure intact. A chemical bond is primarily formed by the redistribution of electrons between two or more atoms of the same or different elements.
Electronic redistribution may occur through the mutual sharing of electrons or by the loss or gain of electrons from one atom’s valence (outermost) shell to another. Atoms develop chemical bonding in order to gain stability, i.e., to achieve a stable noble gas electronic configuration.
The energy changes taking place during a chemical reaction largely dominate the breaking of an old bond (on the reactant side) and/or the formation of a new chemical bond (on the product side).
Rules and energetics of chemical bond formation
Chemical bond breaking requires energy. Contrarily, the formation of a new chemical bond releases energy. If the amount of energy released during bond formation is greater than the amount of energy absorbed during bond breaking, then the chemical reaction is termed exothermic.
A negative enthalpy change (-∆H) symbolizes an exothermic reaction. A positive enthalpy change (+∆H) indicates an endothermic chemical reaction in which the amount of energy absorbed is greater than the amount of energy released. The lower the potential energy of a chemical bond, the more stable it is.
When two atoms approach each other, the positively charged nucleus of one atom attracts the negatively charged electrons of the other and vice versa. Conversely, the same charge electrons of the two atoms experience a strong repulsive effect. The system’s potential energy decreases with the decrease in distance between the two atoms. At the internuclear distance where the force of attraction overcomes the force of repulsion, a chemical bond is formed.
The energy of the system is lowest at this point; the chemical bond is the most stable here. Thus, it is known as the bond energy. The internuclear distance i.e., the distance two atomic nuclei at the point of chemical bond formation is known as the bond length.
Further decrease in distance between the two nuclei leads to a stronger repulsive effect; thus, the chemical bond breaks, and the energy of the system excites again.
Main types of chemical bonds
The following main types of chemical bonds are formed in chemistry.
An ionic bond, also known as an electrovalent bond, is formed by the complete transference of electrons.
A positively charged ion is formed by the loss of valence electrons from a neutral atom. It is known as a cation. Contrarily, a negatively charged ion is formed by the gain of electrons. It is known as an anion. As the electrons transfer from a cation to an anion, a chemical force of attraction develops between the oppositely charged ions. It is known as an ionic bond.
An ionic bond is usually formed between a metal and a non-metal. The metal atom predominantly comes from Group I A (alkali metals) or Group II A (alkaline earth metals) of the Periodic Table of elements. However, the non-metals atoms are from Group VI A or Group VII A (halogens) of the Periodic Table.
Example: Ionic bonding between sodium (Na) and chlorine (Cl2).
- Sodium (Na) metal is present in Group I A. It has only a single electron in its valence shell. In order to gain a stable octet electronic configuration; Na can readily lose this electron and form a monovalent Na+ cation.
- Chlorine (Cl2) is a halogen from Group VII A. It has 7 valence electrons. It is thus short of a single electron in order to complete its octet. So the Cl atom gains 1 electron lost by Na+.
- An electrostatic force of attraction develops between the oppositely charged Na+ and Cl– ions, and an ionic bond is formed.
- A large number of NaCl unit cells arranged as a three-dimensional crystal lattice produce NaCl ionic crystal. This gives us the common table salt, i.e., sodium chloride (NaCl).
- Extremely strong forces of attraction exist in the ionic crystal; therefore, a tremendous amount of energy, called lattice energy, is required to separate constituent NaCl unit cells. Consequently, sodium chloride has a high melting point (801°C).
A covalent chemical bond, also known as an electron pair bond, is formed by the mutual sharing of electrons between two atoms.
More than one covalent chemical bond can be formed between the concerning atoms, such as a double covalent bond and/or a triple covalent bond. In a triple covalent bond, the first bond formed between two atoms is always a sigma (σ) bond. The second and third bonds are pi (π) bonds.
A sigma bond is formed by the overlap of atomic orbitals in such a way that the electron cloud stays uniformly distributed along the nuclear axis. A comparatively weaker pi bond is formed by an orbital overlap in which the shared electron cloud lies above and below the nuclear axis.
A covalently bonded molecule could either be homoatomic or heteroatomic.
- A homoatomic molecule is made up of atoms of the same element for, e.g., H2, Cl2, O3, etc.
- A heteroatomic molecule is made up of atoms from two or more different elements, e.g., H2O, CO2, NH3, H2SO4, etc.
A covalently bonded molecule could also be monoatomic, diatomic, triatomic, polyatomic etc.
- The monoatomic molecule is a special case. Its examples include noble gases such as He, Ne, Ar, etc. Under ordinary conditions, the noble gas atoms do not react with other atoms of the same or different elements. This is because they are already stable; thus, they stay inert. So, in accordance with this, a monoatomic molecule is made up of only a single atom. In reality, no covalent bonding exists in monoatomic molecules, so these are exceptions.
- A diatomic molecule is composed of two atoms of the same or different elements. For e.g., H2 is a homoatomic diatomic molecule. CO is a heteroatomic diatomic molecule.
- A triatomic molecule comprises three same or different types of atoms. O3 is a homoatomic triatomic molecule. SO2 is a heteroatomic triatomic molecule.
- A polyatomic molecule comprises more than three atoms from different elements such as H2SO4.
Covalently bonded molecules could also be neutral such as H2SO4, and/or charged, such as the sulfate [SO4]2- ion is a negatively charged molecular ion.
Similarly, both polar and non-polar covalently bonded molecules can be formed. H2O is a polar, covalently bonded molecule, while N2, CO2, CH4, etc., are non-polar in nature. The polarity of a covalent bond depends on the electronegativity difference between the bonded atoms.
A pure polar covalent bond consists of bonded atoms having an electronegativity difference between 0.5 to 1.6 units. At an electronegativity difference greater than 1.6 units, the covalent bond starts displaying ionic characteristics. In purely ionic bonds, the bonded atoms have an electronegativity difference greater than 2.0 units.
A large number of covalently bonded molecules held together in a giant molecular lattice forms solids such as diamond and graphite.
In simple molecules such as water, covalently bonded molecules are held together via intermolecular forces of attraction.
There are four main types of intermolecular forces of attraction:
- Van der Waal’s forces: Distance-dependent electric forces that keep neutral molecules together.
- London dispersion forces: The intermolecular force of attraction that develops by the induction of temporary dipoles in neutral molecules. It is a part of the Vander Waal’s forces.
- Dipole-dipole forces: The force of attraction between permanently charged molecules.
- Hydrogen bonding: The intermolecular force of attraction between the partial positively charged hydrogen of one molecule to the partial negatively charged electronegative atom (such as O or N) of an adjacent molecule.
H-bonding is the strongest intermolecular force of attraction out of all those mentioned above. However, we must remember that an intermolecular force of attraction between the molecules is always weaker than a much stronger intramolecular chemical bonding, such as covalent bonding within the respective molecules.
Difference between an ionic and a covalent bond
|Complete transference of electrons
|Mutual sharing of electrons
|Usually stronger than a covalent bond
|Usually weaker than an ionic bond
|Ionic compounds are usually solids at r.t.p.
|Covalently bonded molecules are usually liquid and/or gases at r.t.p.
|Chemical bonding between metals and non-metal atoms
|Chemical bonding between non-metal atoms
|Ionic compounds are good conductors of heat and electricity in molten or aqueous forms.
|Covalently bonded molecules are largely poor conductors of heat and electricity in all forms.
Coordinate covalent bond
A coordinate covalent bond, also known as a dative covalent bond, is a type of covalent chemical bond in which the electrons are fully shared by one specie, rather than mutual sharing of electrons. The atom that shares its electrons with another atom is known as the donor, while the other is an acceptor.
Example: Coordinate covalent bonding between NH3 and H+.
In the ammonia (NH3) molecule, a lone pair of electrons is present on the central nitrogen (N) atom. The nitrogen donates this lone pair to form a covalent chemical bond with the electron-deficient hydrogen [H]+ ion. Nitrogen is the donor atom in this example, while H+ is the acceptor. The arrow (see the figure below) points from the donor to the acceptor specie. The ammonia molecule acts as a Lewis base, while H+ acts as a Lewis acid. An ammonium [NH4] + ion is formed as a result of this chemical bonding.
Once the coordinate covalent bond is formed, it is treated as any other ordinary covalent bond in terms of reactivity and properties.
A metallic bond refers to the electrical force of attraction between positive metal ions and the delocalized electrons in a metal structure.
Metals are generically electropositive in nature. This means the metal atoms tend to lose electrons in order to achieve a stable noble gas electronic configuration. In a metallic structure, such as that of the lithium (Li) metal, the Li atoms lose electrons to produce Li+ ions. As a large number of electrons are lost, numerous Li+ ions are simultaneously produced. This results in a sea of delocalized electrons containing positively charged ions.
The strong force of attraction that develops between the positively charged ions and the negatively charged electrons is thus known as metallic bonding.
Different theories of chemical bonding
Different chemical bond theories proposed by scientists help us in predicting the structure and properties related to a specific chemical bond.
Let’s see what these different theories of chemical bonding are.
The valence shell electron pair repulsion (VSEPR) theory of chemical bonding predicts the shape and geometry of a covalently bonded molecule.
According to the VSEPR theory, the molecule adopts a shape such that the electronic repulsions are minimized between the bonded atoms and the lone pairs surrounding a central atom.
VBT stands for the valence bond theory. It discusses the process of hybridization during chemical bond formation.
According to VBT , during chemical bonding different energy atomic orbitals combine to form the same energy hybrid orbitals. The hybrid orbitals then overlap to form new chemical bonds.
The molecular orbital theory (MOT) states that once chemical bonds are formed, the individual atomic orbitals combine to form molecular orbitals.
In this way, the shared electrons are not localized between their parent atomic nuclei. Rather, the electrons belong to the whole molecule containing multiple nuclear centers.
The linear combination of atomic orbitals in the same phase yields a low-energy bonding molecular orbital. Contrarily, different phase molecular orbitals combine to form the high-energy molecular orbitals called anti-bonding molecular orbitals. The shared electrons fill these molecular orbitals in an ascending energy order.
Electron gas theory
The electron gas theory, also known as electron sea theory, explains metallic bonding. It states that positively charged metal ions do not float randomly in the sea of delocalized electrons. Rather, they occupy definitely measurable positions.
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2. House, James E., and Kathleen A. House. 2016. ‘Chapter 4 – Ionic Bonding, Crystals, and Intermolecular Forces.’ in James E. House and Kathleen A. House (eds.), Descriptive Inorganic Chemistry (Third Edition) (Academic Press: Boston).
3. Moeller, Therald, John C. Bailar, Jacob Kleinberg, Cyrus O. Guss, Mary E. Castellion, and Clyde Metz. 1980. ‘9 – THE CHEMICAL BOND.’ in Therald Moeller, John C. Bailar, Jacob Kleinberg, Cyrus O. Guss, Mary E. Castellion and Clyde Metz (eds.), Chemistry (Academic Press).
4. Ouellette, Robert J., and J. David Rawn. 2014. ‘1 – Structure and Bonding in Organic Compounds.’ in Robert J. Ouellette and J. David Rawn (eds.), Organic Chemistry (Elsevier: Boston).
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