Ionic vs Covalent bonds; Ionic bonding requires the attraction of oppositely charged ions, however, in covalent bonding, a stability is acquired through sharing electrons. The configuration and shape is determined by valence shell electron pair repulsion VSEPR rules. In ionic bonding, the geometry is determined by packaging rules and there are no shared electron pairs. The chemistry and properties defining each bond type is detailed below.

Ionic vs covalent bonds (which are stronger)?

An ionic bond allows an amount of sharing of electron density. It is defined when this is larger than the covalent character. The bigger the difference in electronegativity, the more ionic. Bonds with a hybrid character are called polar covalent bonds. Pauling estimated that an electronegativity difference of 1.7 corresponds to 50% ionic character, so that a difference greater than 1.7 corresponds to a bond which is ionic.

Quadrupolar nuclei

Ionic character in covalent bonds can be directly measured for atoms having quadrupolar nuclei (²H, ¹⁴N, ^81,79Br, ^35,37Cl). Interactions between the nuclear quadrupole moments Q and the electric field gradients (EFG) are characterised via the nuclear quadrupole coupling constants:

QCC = e²q_zzQ/h

where the eq_zz term corresponds to the principal component of the EFG tensor and e is the elementary charge.

As ionic bonding occurs, it is inaccurate to define a single ionic bond between two atoms, as the binding forces that keep the lattice stable are more of a collective force. In covalent bonding, there is a definitive, localised bond between two atoms.

What is a covalent bond? (definition and properties)

A covalent bond is a chemical bond that involves the sharing of electron pairs. These bonding pairs, the stable balance of attractive and repulsive forces between atoms, are known as covalent bonds. This allows each atom to attain the equivalent of a full outer shell. Covalent bonding also includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, bent bonds, three-center two-electron bonds and three-center four-electron bonds.

The meaning of covalent bonding

The term covalent bond dates from 1939. The prefix co- means jointly, thus a co-valent bond, means that the atoms share valence. Hydrogen atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities.

Types of covalent bonds

The term covalence was first used in 1919 by Langmuir and he wrote that “we shall denote by the term covalence the number of pairs of electrons that a given atom shares with its neighbours.”

History of covalent bonding

Lewis, in 1916 described the sharing of electron pairs. The Lewis notation, where valence electrons are represented as pairs of dots between atoms. Atom must form enough covalent bonds to form a full outer electron shell. For methane, the carbon atom has a valence of four and is surrounded by eight electrons, four from the carbon itself and four from the hydrogens. Each hydrogen has a valence of one and is surrounded by two electrons, which is called a duet rule.

Quantum mechanics accurately predicts the nature of the bonds and the structures and properties of molecules. Heitler and London are credited with the first successful quantum mechanical explanation of a chemical bond in 1927.

Pi and Sigma atomic bonds

Atomic orbitals (except s orbitals) have directional properties leading to different types of bonds. Sigma (σ) bonds are the strongest covalent bonds and are due to head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond. Pi (π) bonds are weaker and are due to lateral overlap between p (or d) orbitals. A double bond between two given atoms consists of one σ and one π bond.

Covalent bonds are also affected by the electronegativity of the connected atoms which determines the chemical polarity of the bond. Two atoms with equal electronegativity will make non-polar covalent bonds such as H–H. An unequal relationship creates a polar covalent bond such as with H−Cl. Polarity also requires geometric asymmetry, or else dipoles may cancel out resulting in a non-polar molecule.

Covalent structures

Different structures exist including individual molecules, macromolecular structures and giant covalent structures. Individual molecules have negligible forces of attraction, e.g. gases HCl, SO₂ and CH₄. In molecular structures, there are also weak forces of attraction. Such covalent substances are low-boiling-temperature liquids, like ethanol, and low-melting-temperature solids, like iodine.

Macromolecular structures

These have large numbers of atoms linked by covalent bonds in chains, including synthetic polymers such as polyethylene and proteins.

Giant covalent structures or network covalent structures

These contain large numbers of atoms linked e.g. graphite. These substances have high melting and boiling points and tend to have high electrical resistivity. Elements that have high electronegativity, and the ability to form three or four electron pair bonds, form large macromolecular structures.

One and three electron bonds

Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the dihydrogen cation, H⁺₂. One-electron bonds often have about half the bond energy of a 2-electron bond. However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li⁺₂ than for the 2-electron Li₂.

The simplest example of three-electron bonding can be found in the helium dimer cation, He⁺₂. It is considered a half bond because it consists of only one shared electron; the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Chlorine dioxide and iodine dioxide also contain three-electron bonds.

Resonance

A single Lewis structure maybe insufficient to explain the electron configuration, hence a superposition of structures is needed. The same two atoms in such molecules can be bonded differently in different structures (a single bond in one, a double bond in another, or even none at all), resulting in a non-integer bond order. The nitrate ion is one such example with three equivalent structures. The bond between the nitrogen and each oxygen is a double bond in one structure and a single bond in the other two, so that the average bond order for each N–O interaction is 2 + 1 + 1/3 = 4/3.

Aromaticity

In organic chemistry, when a molecule with a planar ring obeys Hückel’s rule, where the number of π electrons fit the formula 4n + 2 (where n is an integer), it gains extra stability and symmetry. In benzene, the prototypical aromatic compound, there are 6 π bonding electrons (n = 1, 4n + 2 = 6). These occupy three delocalised π molecular orbitals or form conjugate π bonds in two resonance structures that combine, creating a hexagon exhibiting a greater stabilisation.

Hypervalence

Xenon difluoride and sulfur hexafluoride have higher co-ordination numbers than expected due to the octet rule. This is the three-center four-electron bond (3c–4e) model which interprets the molecular wavefunction in terms of non-bonding highest occupied molecular orbitals in molecular orbital theory and resonance of sigma bonds in valence bond theory.

Quantum mechanical description valence chemical bonding

After the development of quantum mechanics, two basic theories were proposed to provide a quantum description of chemical bonding: valence bond (VB) theory and molecular orbital (MO) theory. A recent quantum description is given in terms of atomic contributions to the electronic density of states.

Comparison of VB and MO theories

These are two ways to build up the electron configuration. For valence bond theory, the atomic hybrid orbitals are filled with electrons first to produce a fully bonded valence configuration, followed by performing a linear combination of contributing structures (resonance). In molecular orbital theory a linear combination of atomic orbitals is followed by filling of the resulting molecular orbitals with electrons.

As valence bond theory builds the molecular wavefunction out of localised bonds, it is more suited for the calculation of bond energies and the appreciation of mechanisms. As molecular orbital theory builds the molecular wavefunction out of delocalised orbitals, it is more suited for the calculation of ionization energies.

Simple (Heitler–London) valence bond theory correctly predicts the dissociation of homonuclear diatomic molecules into separate atoms, while Hartree–Fock molecular orbital theory incorrectly predicts dissociation into a mixture of atoms and ions. However, simple molecular orbital theory correctly predicts Hückel’s rule of aromaticity, while simple valence bond theory incorrectly predicts that cyclobutadiene has larger resonance energy than benzene.

Although the wavefunctions do not agree and do not match the stabilisation energy by experiment, they can be corrected by configuration interaction.  This is done by combining the valence bond covalent function with the functions describing all possible ionic structures or by combining the molecular orbital ground state function with the functions describing all possible excited states using unoccupied orbitals. It can then be seen that the simple molecular orbital approach overestimates the weight of the ionic structures while the simple valence bond approach neglects them.

Modern calculations in quantum chemistry usually start from a molecular orbital rather than a valence bond approaches the MO approach is more readily adapted to numerical computations. Molecular orbitals are orthogonal, which increases the feasibility of computer calculations. However, better valence bond programs are becoming available.

Ionic bonding definition

Atoms that have an almost full or almost empty valence shell are very reactive. Atoms that are strongly electronegative frequently bond with other molecules or gain electrons to form anions. Atoms that are weakly electronegative (such as alkali metals) have relatively few valence electrons and can be shared with atoms that are strongly electronegative. As a result, weakly electronegative atoms tend to distort their electron cloud and form cations.

Ionic bond formation

Ionic bonds output from a redox reaction when atoms of a metallic element, whose ionisation energy is low, offer some of their electrons to achieve a stable configuration, forming cations. An atom of another element, usually non-metal with greater electron affinity accepts one or more electrons, forming anions. The electrostatic attraction creates the formation of a solid crystallographic lattice.

For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na⁺), and the chlorine atoms each gain an electron to form anions (Cl⁻).

Na + Cl → Na⁺ + Cl⁻ → NaCl

However, to maintain neutrality, strict ratios between anions and cations ensure stoichiometry rules are followed. For compounds that possess mixed ionic and metallic bonding, this may not be true. Many sulfides do form non-stoichiometric compounds. Many ionic compounds are referred to as salts as they can also be formed by the neutralisation reaction of an Arrhenius base like NaOH with an Arrhenius acid like HCl.

NaOH + HCl → NaCl + H₂O

The removal of electrons to form the cation is endothermic, raising the system’s energy. There may also be energy changes associated with breaking of bonds. The anion accepting the cation’s valence electrons and the subsequent attraction of the ions releases lattice energy and lowers the overall energy of the system.

Ionic compounds

Ionic compounds in the solid state form lattice structures. The two factors determining the form of the lattice are the relative charges of the ions and their sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as magnesium oxide.

Strength of ionic bonding

The enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born–Haber cycle. The electrostatic potential can be expressed in terms of the inter-ionic separation and the Madelung constant, which considers the geometry. The Born-Landé equation gives a reasonable fit to the lattice energy e.g. sodium chloride, where the calculated value is −756 kJ/mol, which compares to −787 kJ/mol using the Born–Haber cycle.

Bjerrum or Fuoss equation

In aqueous solution the binding strength can be described by the Bjerrum or Fuoss equation as function of the ion charges, independent of the nature of the ions such as polarisability or size. The strength of salt bridges is calculated by measurements of equilibria between molecules containing cationic and anionic sites. Equilibrium constants in water indicate additive free energy contributions for each salt bridge. The attractive forces can be modelled by Coulomb’s Law.

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