9 classical theories of chemical bonding

This article is an overview of the different theories of chemical bonding. Scientists proposed various models and theories over time in order to explain how a new chemical bond is formed. These included classical theories, the age-old ideas that remained relevant till today to modern theories, most applicable to the idea of a chemical bond in the 21^st century.

Continue reading the article to learn about these different chemical bonding theories. But let’s first start with a brief introduction to chemical bonding and different types of chemical bonds.

What is chemical bonding?

Chemical bonding lies at the heart of all sciences and every substance that we have ever known. It is defined as a force of attraction that holds atoms or ions together in a stable, independent entity such as a molecule or a crystal.

A chemical bond is primarily formed by the redistribution of electrons between two or more atoms of the same or different elements.

Atoms combine in a molecule by developing chemical bonding in order to decrease their overall potential energy and thus gain stability.

Different types of chemical bonds

There are three main types of chemical bonds:

i) Ionic bond: An ionic bond is formed by the complete transference of electrons from one atom to another, such as a metal to a non-metal atom. A large number of ionic units held together in a three-dimensional arrangement forms an ionic lattice.

ii) Covalent bond: A covalent bond is formed by the mutual sharing of electrons between two non-metal atoms. A coordinate covalent bond or dative covalent bond is a sub-type of the covalent chemical bond. It is formed when both the electrons shared are provided by a donor atom while the other constituent atom acts as an acceptor.

iii) Metallic bond: Metallic bonding develops between different metal atoms in a metallic crystal. Metals are electropositive species that lose electrons and attain positive charges. The force of attraction between positively charged cations and negative electrons in a sea of delocalized electrons is called a metallic bond.

To explain chemical bonding in the above three different types of bonds, a series of bonding theories have been introduced over time. Let us discuss each, one by one.

1. Electronic theory

• Electronic theory mainly focuses on explaining chemical bonding in ionic compounds.
• An ionic bond is also known as an electrovalent bond.
• As per the electronic theory, neutral atoms change into ions as they tend to achieve the stable electronic configuration of their nearest noble gas element.
• For instance, sodium (Na) has one excess electron in its valence shell. So, it loses this single electron to transform into Na⁺ ion (cation). It thus achieves a complete octet and an electronic configuration (1s² 2s² 2p⁶) identical to Neon (¹⁰Ne).

• Chlorine (Cl) has a deficiency of one electron. So, it gains the electron lost by Na and transforms into Cl^– (anion) by achieving a complete octet (1s² 2s² 2p⁶ 3s² 3p⁶), i.e., the electronic configuration of Argon (¹⁸Ar).

• Electrostatic forces of attraction develop between oppositely charged Na⁺ and Cl^– ions to form NaCl.

• A large number of [Na⁺ Cl^–] ion pairs arrange three-dimensionally to form a NaCl crystal lattice. Each Na⁺ ion is surrounded by 6 Cl^– ions and vice versa in the cubic lattice arrangement.

2. Lewis bonding theory

• A Lewis dot structure is a simplified representation of all the valence electrons present in a chemically bonded molecule.

• As per Lewis bonding theory, only the valence (outermost) shell electrons of an atom get involved in chemical bonding.
• A new chemical bond is formed by complete transference (ionic bond) or mutual sharing (covalent bond) of the valence electrons from constituent atoms.
• In this way, the Lewis model is applicable for representing both ionic as well as covalently bonded molecules (as shown below).

• As per Lewis theory, the lower the formal charges present on the bonded atoms in a covalent molecule, the higher its stability.
• Lewis models also predict polarity and resonance present in different covalently bonded molecules.

3. Electron sea theory

• The electron sea theory is a simple approach to explain metallic bonding. It is also known as the electron gas theory.
• As per electron sea theory, metal cations are embedded in a sea of delocalized electrons.
• The electrons are free to move throughout the structure. However, metal cations occupy fixed positions at measurable distances from each other.
• The opposite charge attraction that develops between metal cations and electrons in this structure is called metallic bonding.

The presence of delocalized electrons, as per the electron sea theory, is extremely beneficial in explaining the versatile properties of metals such as their:

i). Ability to conduct heat and electricity.

ii). Metallic lustre.

iii). Malleability and ductility.

iv). Elasticity.

v). Photoelectric effect.

4. Valence shell electron pair repulsion (VSEPR) theory

• The VSEPR theory is predominantly used for explaining covalent bonding.
• The electron pairs present on a covalently bonded molecule can be distinguished into bond pairs and lone pairs (non-bonded electron pairs) using their Lewis structures.
• As per the VSEPR theory, these electron pairs arrange themselves in a molecule such that the overall electron repulsive effect is minimized.
• In this way, the VSEPR theory predicts the shape and geometry of covalent molecules.
• The shapes of molecules can be predicted using the AXE formula as per the VSEPR concept.

• A stands for the central atom in a polyatomic molecule. X represents the bonded atoms, while E denotes the lone pairs of electrons present on the central atom in a molecule.
• The electron repulsive effect increases in the order: bond pair-bond pair repulsions > bond pair-lone pair repulsions > lone pair-lone pair repulsions.
• Therefore, as bond pairs get replaced by lone pairs, the electron repulsive effect increases, which leads to distortion in the shape and geometry of the molecule. The X-A-X bond angle decreases.

Refer to the table given below in order to see how molecules adopt different shapes and geometry as per the AXN formula.

5. Valence bond theory (VBT)

• The valence bond theory (VBT) introduced the idea of orbital overlap while covalent bond formation.
• As per VBT, different energy atomic orbitals mix to generate the same energy hybrid orbitals during chemical bonding. This concept is known as orbital hybridization.
• The formation of hybrid orbitals allows sufficient orbital overlap in order to form a new chemical bond.
• VBT assumed that bonded electrons stay localized or fixed in the overlapping zone.
• The atomic orbitals regain their individuality post-chemical bonding.

• The quantum mechanical treatment of VBT gave the first quantitative description of a covalent chemical bond.
• This concept was further extended to discuss coordination complexes and metallic bonding.
• As per resonance VBT, a solid metallic structure is believed to involve electron pair resonance between each atom and its neighbouring atoms. In this way, VBT proposed the existence of covalent bonds in metallic structures as well. The resonance model of lithium metal is shown below.

• The valence bond theory was also applied to explain the bonding present in coordination complexes. A coordination complex consists of a metal atom or ion surrounded by electron donors called ligands. As per VBT of coordination complexes, different energy atomic orbitals of the central metal hybridize to produce the same energy hybrid orbitals. The empty hybrid orbitals accommodate extra electrons from the ligands in a metal complex.

• For instance, [CoCl₄]^2- is a tetrahedral coordination complex. In this, Co²⁺ has sp³ hybridization. An electron pair from each Cl^– ion is accommodated in an empty sp³ hybrid orbital of Co²⁺. In contrast, [Ni(CN)₄]^2- is a square planar complex. The central Ni²⁺ ion is dsp² hybridized in this case.

• In both the above examples, the coordination number of the central metal ion is 4.
• However, the hybridization and shape of the complex are governed by the strength of incoming ligands as per VBT. 

6. Molecular orbital theory (MOT)

• VBT explains what happens during bond formation, while MOT is an after-bonding phenomenon to explain covalently bonded molecules.
• As per MOT, individual atomic orbitals combine to form new molecular orbitals (MOs) after chemical bonding.
• The linear combination of ‘n’ atomic orbitals (n = integer) produces an equal number of molecular orbitals.
• For example, 2 atomic orbitals of hydrogen atoms in H₂ combine to form 2 MOs, including a bonding MO and an anti-bonding MO.

• As per the quantum mechanical treatment, bonding MO is formed by constructive interference (additive effect) of two wave functions. However, an antibonding MO is formed by their destructive interference (subtractive effect). 

• The bonded electrons present in the molecule thus belong to all nuclear centers and are not confined in a specific overlapping region.

• In this way, MOT is an advanced bonding theory that negates VBT’s idea of localized electrons in a covalently bonded molecule.

7. Band theory

• The band theory is also sometimes known as the molecular orbital theory of metals.
• As per the band theory, when a huge number of metal atoms interact in a metallic crystal, their atomic orbitals combine to form molecular orbitals that belong to the crystal as a whole.
• An Avogadro’s number (6.02 x 10²³) of atomic orbitals combine to form an Avogadro’s number of molecular orbitals with a minute energy difference.
• It is due to this small energy difference that the molecular orbitals appear to be a continuum called a quasi-continuous energy band.
• Band theory finds its utmost relevance in determining the electrical conductivity of metals.
• Delocalization of electrons is an essential condition for the conduction of electrical current.
• The energy band formed by the combination of inner shell atomic orbitals such as 1s, 2s and 2p in sodium metal is completely filled with electrons. Hence it is known as a non-conduction band.
• The energy band formed by 3s orbitals in half-filled. It is the valence band.
• The energy band formed by 3p orbitals is empty.
• On providing energy, the 3s electrons overcome the band gap.
• The 3s valence electrons thus easily jump to the empty 3p band where they are free to move and thus conductivity electricity. So, the 3p band is known as the conduction band in sodium metal.

 8. Crystal field theory (CFT)

• The crystal field theory (CFT) explains chemical bonding in transition metal complexes.
• As per CFT, purely electrostatic forces of attraction develop as ligands (anions or neutral molecules) approach a metal atom or ion in a transition metal complex.
• The ligands behave as point charges or point dipoles.
• The electric field of the ligands repels the electrons present in the d-orbitals of the central metal.
• The five degenerate d-orbitals thus undergo crystal field splitting by dividing into two groups lying at different energy levels.
• In an octahedral metal complex, d_z² and d_x²_-y² form the higher energy set called the, e_g set. In contrast, the d_xy, d_yz and d_xz occupy a lower energy level known as the t_2g set.  
• The energy difference between the e_g set and t_2g set of orbitals is known as crystal field splitting energy or crystal field stabilization energy of the complex.
• It is calculated by applying the formula: Crystal field stabilization energy ∆_o (octa) = -0.4 (m) t_2g + 0.6 (n) e_g where m and n are the electrons present in t_2g and e_g sets, respectively.
• The electrons are filled in these different energy orbitals following Aufbau, Hund’s rule and Pauli Exclusion Principle.

• The d electrons get paired up in lower energy orbitals before occupying the higher energy set if the ligand is a strong field ligand.
• In that case, the spin-pairing energy is lower than the crystal field splitting energy ∆_o. This results in a spin-paired or low-spin diamagnetic complex.
• Contrarily, a high spin paramagnetic complex is formed in the presence of a weak field ligand, in which case the spin pairing energy is greater than ∆_o.
• Stronger the ligand strength, the higher the ∆_o value; thus more stable the complex formed.
• CFT is quite useful in explaining the bright colors exhibited by metal complexes in addition to predicting their stability, reactivity and magnetic behaviour. 

9. Ligand field theory (LFT)

• The ligand field theory (LFT) is also known as the molecular orbital theory (MOT) of transition metal complexes.
• It is by far the most modern bonding theory to explain chemical bonding in metal complexes.
• LFT is more advanced than CFT as it incorporates the role of s and p orbitals as well during metal complex formation, while CFT is only focused on explaining the behavior of d-orbitals.
• As per LFT, when a transition metal complex is formed, s, p and d orbitals of the central metal atom or ion that are in right symmetry with the ligand orbitals undergo a linear combination of atomic orbitals (LCAO).
• This results in the formation of new orbitals called molecular orbitals, including bonding and non-bonding molecular orbitals.
• The metal orbitals, which are off-symmetry with ligand orbitals, stay non-bonded.
• Electrons are placed in all the above orbitals following Aufbau, Hund’s rule and Pauli Exclusion principles, strictly influenced by ligand strength.
• Following the placement of electrons in the MO diagram (as shown for [Co(NH₃)₆]³⁺), we can predict the stability, magnetism, color and reactivity of metal complexes.

• LFT also explains the formation of highly stable metal complexes in the presence of pi-acceptor ligands, following the concept of donation followed by back donation of electrons from metal (M) to ligand (L) in an M-L interaction.  

All the images and diagrams provided in this article are designed by the writer herself (Ammara W.)

References

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3. H. Crabtree, Robert. 2014. The organometallic chemistry of the transition metals (Wiley).

4. Neese, Frank. 2013. ‘Chapter 2 – Introduction to Ligand Field Theory.’ in Robert R. Crichton and Ricardo O. Louro (eds.), Practical Approaches to Biological Inorganic Chemistry (Elsevier: Oxford).

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