Chemical bonding involving attraction between ions
Sodium andfluorine atoms undergoing a redox reaction to form sodium ions and fluoride ions. Sodium loses its outerelectron to give it a stableelectron configuration, and this electron enters the fluorine atomexothermically. The oppositely charged ions – typically a great many of them – are then attracted to each other to form solidsodium fluoride.
Ionic bonding is a type ofchemical bonding that involves theelectrostatic attraction between oppositely chargedions, or between twoatoms with sharply differentelectronegativities,[1] and is the primary interaction occurring inionic compounds. It is one of the main types of bonding, along withcovalent bonding andmetallic bonding. Ions are atoms (or groups of atoms) with an electrostatic charge. Atoms that gain electrons make negatively charged ions (calledanions). Atoms that lose electrons make positively charged ions (calledcations). This transfer of electrons is known aselectrovalence in contrast tocovalence. In the simplest case, the cation is ametal atom and the anion is anonmetal atom, but these ions can be more complex, e.g.polyatomic ions likeNH+ 4 orSO2− 4. In simpler words, an ionic bond results from the transfer of electrons from ametal to anon-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 ofcovalent bonding or electron sharing. Thus, the term "ionic bonding" is given when the ionic character is greater than the covalent character – that is, a bond in which there is a large difference inelectronegativity between the cation and anion, causing the bonding to be more polar (ionic) than in covalent bonding where electrons are shared more equally. Bonds with partially ionic and partially covalent characters are calledpolar covalent bonds.[2]
Ionic compounds conductelectricity when molten or in solution, typically not when solid. Ionic compounds generally have a highmelting 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 besoluble in water; the stronger the cohesive forces, the lower the solubility.[3]
Overview
Atoms that have an almost full or almost emptyvalence shell tend to be veryreactive. Strongly electronegative atoms (such ashalogens) often have only one or two empty electron states in theirvalence shell, and frequentlybond with other atoms or gain electrons to formanions. Weakly electronegative atoms (such asalkali metals) have relatively fewvalence electrons, which can easily be lost to strongly electronegative atoms. As a result, weakly electronegative atoms tend to distort theirelectron cloud and formcations.
Properties of ionic bonds
They are considered to be among thestrongest of all types of chemical bonds. This often causes ionic compounds to be very stable.
Ionic bonds havehighbond 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 acrystal structure, in which the ions occupy the corners of the crystal. Such a structure is called acrystal lattice.
Ionic compoundslose their crystal lattice structure and break up into ions when dissolved inwater or any otherpolar 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 theirmelting point in a process known asmelting.
Formation
Ionic bonding can result from aredox reaction when atoms of an element (usuallymetal), whoseionization energy is low, give some of their electrons to achieve a stable electron configuration. In doing so, cations are formed. An atom of another element (usually nonmetal) with greaterelectron affinity accepts one or more electrons to attain a stableelectron configuration, and after accepting electrons an atom becomes an anion. Typically, the stable electron configuration is one of thenoble gases for elements in thes-block and thep-block, and particularstable electron configurations ford-block andf-block elements. The electrostatic attraction between the anions and cations leads to the formation of a solid with acrystallographic 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 betweenlithium andfluorine to formlithium fluoride. Lithium has a low ionization energy and readily gives up its lonevalence 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 isisoelectronic withhelium and fluorine is isoelectronic withneon. 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, commontable salt issodium chloride. Whensodium (Na) andchlorine (Cl) are combined, the sodium atoms each lose anelectron, 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 assalts 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 toCoulomb'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.
In the rock salt lattice, each sodium ion (purple sphere) has anelectrostatic 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 saltsodium chloride is also adopted by manyalkali halides, and binary oxides such asmagnesium oxide.Pauling's rules provide guidelines for predicting and rationalizing the crystal structures of ionic crystals
For a solid crystalline ionic compound theenthalpy change in forming the solid from gaseous ions is termed thelattice energy. The experimental value for thelattice energy can be determined using theBorn–Haber cycle. It can also be calculated (predicted) using theBorn–Landé equation as the sum of theelectrostatic potential energy, calculated by summing interactions between cations and anions, and a short-range repulsive potential energy term. Theelectrostatic 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. TheBorn–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 theBorn–Haber cycle.[4][5] In aqueous solution the binding strength can be described by theBjerrum 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 in complicated molecules iscrystallography, sometimes also NMR-spectroscopy.
The attractive forces defining the strength of ionic bonding can be modeled byCoulomb's Law. Ionic bond strengths are typically (cited ranges vary) between 170 and 1500 kJ/mol.[8][9]
Polarization power effects
Ions incrystal lattices of purely ionic compounds arespherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised inFajans' rules. Thispolarization of the negative ion leads to a build-up of extra charge density between the twonuclei, that is, to partial covalency. Larger negative ions are more easily polarized, but the effect is usually important only when positive ions withcharges of 3+ (e.g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present). Note that this is not theionic polarization effect that refers to the 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 the attraction of oppositely charged ions, whereas, incovalent bonding, atoms are bound by sharing electrons to attain stable electron configurations. In covalent bonding, themolecular geometry around each atom is determined by valence shell electron pair repulsionVSEPR rules, whereas, in ionic materials, the geometry follows maximumpacking rules. One could say that covalent bonding is moredirectional 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 highercoordination 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 sharingelectron density between them. Therefore, all ionic bonding has some covalent character. Thus, bonding is considered ionic where the ionic character is greater than the covalent character. The larger the difference inelectronegativity between the two types of atoms involved in the bonding, the more ionic (polar) it is. Bonds with partially ionic and partially covalent character are calledpolar 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 thePauling 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 or127I). These nuclei are generally objects of NQRnuclear quadrupole resonance and NMRnuclear magnetic resonance studies. Interactions between the nuclear quadrupole momentsQ and the electric field gradients (EFG) are characterized via the nuclear quadrupole coupling constants
QCC =e2qzzQ/h
where theeqzz term corresponds to the principal component of the EFG tensor ande 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 isnot 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 asemiconductor, asemimetal or eventually a metallic conductor with metallic bonding.
^Schneider, Hans-Jörg (2012). "Ionic Interactions in Supramolecular Complexes".Ionic Interactions in Natural and Synthetic Macromolecules. pp. 35–47.doi:10.1002/9781118165850.ch2.ISBN9781118165850.
^David Arthur Johnson,Metals and Chemical Change, Open University, Royal Society of Chemistry, 2002,ISBN0-85404-665-8
^Linus Pauling,The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry, Cornell University Press, 1960ISBN0-801-40333-2doi:10.1021/ja01355a027
^Schneider, H.-J.; Yatsimirsky, A. (2000)Principles and Methods in Supramolecular Chemistry. WileyISBN9780471972532
^Soboyejo, W.O (2003). Mechanical properties of engineered materials. Marcel Dekker. pp. 16–17.ISBN0-203-91039-7.OCLC54091550.
^Askeland, Donald R. (January 2015). The science and engineering of materials. Wright, Wendelin J. (Seventh ed.). Boston, MA. pp. 38.ISBN978-1-305-07676-1.OCLC903959750.
^L. PaulingThe Nature of the Chemical Bond (3rd ed., Oxford University Press 1960) p.98-100.
