Ionic bonding

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exothermically. The oppositely charged ions – typically a great many of them – are then attracted to each other to form solid sodium fluoride
.

Ionic bonding is a type of

non-metal
to obtain a full valence shell for both atoms.

Clean ionic bonding — in which one atom or molecule completely transfers an electron to another — cannot exist: all ionic compounds have some degree of

polar covalent bonds.[2]

Ionic compounds conduct electricity when molten or in solution, typically not when solid. Ionic compounds generally have a high melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the cohesive forces and the higher the melting point. They also tend to be soluble in water; the stronger the cohesive forces, the lower the solubility.[3]

Overview

Atoms that have an almost full or almost empty

cations
.

Properties of ionic bonds

  • They are considered to be among the strongest of all types of chemical bonds. This often causes ionic compounds to be very stable.
  • Ionic bonds have high bond energy. Bond energy is the mean amount of energy required to break the bond in the gaseous state.
  • Most ionic compounds exist in the form of a crystal structure, in which the ions occupy the corners of the crystal. Such a structure is called a crystal lattice.
  • Ionic compounds lose their crystal lattice structure and break up into ions when dissolved in
    polar solvent. This process is called solvation. The presence of these free ions makes aqueous ionic compound solutions good conductors of electricity. The same occurs when the compounds are heated above their melting point in a process known as melting
    .

Formation

Ionic bonding can result from a

crystallographic lattice
in which the ions are stacked in an alternating fashion. In such a lattice, it is usually not possible to distinguish discrete molecular units, so that the compounds formed are not molecular. However, the ions themselves can be complex and form molecular ions like the acetate anion or the ammonium cation.

Representation of ionic bonding between lithium and fluorine to form lithium fluoride. Lithium has a low ionization energy and readily gives up its lone valence electron to a fluorine atom, which has a positive electron affinity and accepts the electron that was donated by the lithium atom. The end-result is that lithium is isoelectronic with helium and fluorine is isoelectronic with neon. Electrostatic interaction occurs between the two resulting ions, but typically aggregation is not limited to two of them. Instead, aggregation into a whole lattice held together by ionic bonding is the result.

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−). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).

Na + Cl → Na+ + Cl− → NaCl

However, to maintain charge neutrality, strict ratios between anions and cations are observed so that ionic compounds, in general, obey the rules of stoichiometry despite not being molecular compounds. For compounds that are transitional to the alloys and possess mixed ionic and metallic bonding, this may not be the case anymore. Many sulfides, e.g., do form non-stoichiometric compounds.

Many ionic compounds are referred to as salts as they can also be formed by the neutralization reaction of an Arrhenius base like NaOH with an Arrhenius acid like HCl

NaOH + HCl → NaCl + H2O

The salt NaCl is then said to consist of the acid rest Cl− and the base rest Na+.

The removal of electrons to form the cation is endothermic, raising the system's overall energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the action of the anion's accepting the cation's valence electrons and the subsequent attraction of the ions to each other releases (lattice) energy and, thus, lowers the overall energy of the system.

Ionic bonding will occur only if the overall energy change for the reaction is favorable. In general, the reaction is exothermic, but, e.g., the formation of mercuric oxide (HgO) is endothermic. The charge of the resulting ions is a major factor in the strength of ionic bonding, e.g. a salt C+A− is held together by electrostatic forces roughly four times weaker than C2+A2− according to Coulomb's law, where C and A represent a generic cation and anion respectively. The sizes of the ions and the particular packing of the lattice are ignored in this rather simplistic argument.

Structures

Main article:
Ionic compound
electrostatic
interaction with its eight nearest-neighbour chloride ions (green spheres)

Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt

alkali halides, and binary oxides such as magnesium oxide. Pauling's rules
provide guidelines for predicting and rationalizing the crystal structures of ionic crystals

Strength of the bonding

Main article: Lattice energy

For a solid crystalline ionic compound the

electrostatic potential can be expressed in terms of the interionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The further away from the nucleus the weaker the shield. The Born–Landé equation gives a reasonable fit to the lattice energy of, e.g., sodium chloride, where the calculated (predicted) value is −756 kJ/mol, which compares to −787 kJ/mol using the Born–Haber cycle.[4][5] In aqueous solution the binding strength can be described by the Bjerrum or Fuoss equation as function of the ion charges, rather independent of the nature of the ions such as polarizability or size.[6] The strength of salt bridges is most often evaluated by measurements of equilibria between molecules containing cationic and anionic sites, most often in solution.[7] Equilibrium constants in water indicate additive free energy contributions for each salt bridge. Another method for the identification of hydrogen bonds also in complicated molecules is crystallography
, sometimes also NMR-spectroscopy.

The attractive forces defining the strength of ionic bonding can be modeled by Coulomb's Law. Ionic bond strengths are typically (cited ranges vary) between 170 and 1500 kJ/mol.[8][9]

Polarization power effects

ionic polarization
effect that refers to displacement of ions in the lattice due to the application of an electric field.

Comparison with covalent bonding

In ionic bonding, the atoms are bound by attraction of oppositely charged ions, whereas, in

packing rules. One could say that covalent bonding is more directional in the sense that the energy penalty for not adhering to the optimum bond angles is large, whereas ionic bonding has no such penalty. There are no shared electron pairs to repel each other, the ions should simply be packed as efficiently as possible. This often leads to much higher coordination numbers
. In NaCl, each ion has 6 bonds and all bond angles are 90°. In CsCl the coordination number is 8. By comparison carbon typically has a maximum of four bonds.

Purely ionic bonding cannot exist, as the proximity of the entities involved in the bonding allows some degree of sharing

polar covalent bonds. For example, Na–Cl and Mg–O interactions have a few percent covalency, while Si–O bonds are usually ~50% ionic and ~50% covalent. Pauling estimated that an electronegativity difference of 1.7 (on the Pauling scale) corresponds to 50% ionic character, so that a difference greater than 1.7 corresponds to a bond which is predominantly ionic.[10]

Ionic character in covalent bonds can be directly measured for atoms having quadrupolar nuclei (2H, 14N, 81,79Br, 35,37Cl or 127I). These nuclei are generally objects of NQR nuclear quadrupole resonance and NMR nuclear magnetic resonance studies. Interactions between the nuclear quadrupole moments Q and the electric field gradients (EFG) are characterized via the nuclear quadrupole coupling constants

QCC = e2qzzQ/h

where the eqzz term corresponds to the principal component of the EFG tensor and e is the elementary charge. In turn, the electric field gradient opens the way to description of bonding modes in molecules when the QCC values are accurately determined by NMR or NQR methods.

In general, when ionic bonding occurs in the solid (or liquid) state, it is not possible to talk about a single "ionic bond" between two individual atoms, because the cohesive forces that keep the lattice together are of a more collective nature. This is quite different in the case of covalent bonding, where we can often speak of a distinct bond localized between two particular atoms. However, even if ionic bonding is combined with some covalency, the result is not necessarily discrete bonds of a localized character.[2] In such cases, the resulting bonding often requires description in terms of a band structure consisting of gigantic molecular orbitals spanning the entire crystal. Thus, the bonding in the solid often retains its collective rather than localized nature. When the difference in electronegativity is decreased, the bonding may then lead to a semiconductor, a semimetal or eventually a metallic conductor with metallic bonding.

See also

References

  1. ISBN 978-0-9678550-9-7
    .
  2. ^ a b Seifert, Vanessa (27 November 2023). "Do bond classifications help or hinder chemistry?". chemistryworld.com. Retrieved 22 January 2024.
  3. ISBN 9781118165850
    .
  4. ISBN 0-85404-665-8
  5. doi:10.1021/ja01355a027
  6. ISBN 9780471972532
  7. PMID 27136957
    .
  8. OCLC 54091550
    .
  9. OCLC 903959750
    .
  10. ^ L. Pauling The Nature of the Chemical Bond (3rd ed., Oxford University Press 1960) p.98-100.

External links

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