In chemistry, the lattice energy is the energy change upon formation of one mole of a crystalline ionic compound from its constituent ions, which are assumed to initially be in the gaseous state. It is a measure of the cohesive forces that bind ionic solids. The size of the lattice energy is connected to many other physical properties including solubility, hardness, and volatility. Since it generally cannot be measured directly, the lattice energy is usually deduced from experimental data via the Born–Haber cycle.[1]

Lattice energy and lattice enthalpy

Sodium chloride crystal lattice

The concept of lattice energy was originally applied to the formation of compounds with structures like rocksalt (NaCl) and sphalerite (ZnS) where the ions occupy high-symmetry crystal lattice sites. In the case of NaCl, lattice energy is the energy change of the reaction

Na+ (g) + Cl (g) → NaCl (s)

which amounts to −786 kJ/mol.[2]

Some chemistry textbooks[3] as well as the widely used CRC Handbook of Chemistry and Physics[4] define lattice energy with the opposite sign, i.e. as the energy required to convert the crystal into infinitely separated gaseous ions in vacuum, an endothermic process. Following this convention, the lattice energy of NaCl would be +786 kJ/mol. Both sign conventions are widely used.

The relationship between the lattice energy and the lattice enthalpy at pressure is given by the following equation:


where is the lattice energy (i.e., the molar internal energy change), is the lattice enthalpy, and the change of molar volume due to the formation of the lattice. Since the molar volume of the solid is much smaller than that of the gases, . The formation of a crystal lattice from ions in vacuum must lower the internal energy due to the net attractive forces involved, and so . The term is positive but is relatively small at low pressures, and so the value of the lattice enthalpy is also negative (and exothermic).

Theoretical treatments

The lattice energy of an ionic compound depends strongly upon the charges of the ions that comprise the solid, which must attract or repel one another via Coulomb's Law. More subtly, the relative and absolute sizes of the ions influence . London dispersion forces also exist between ions and contribute to the lattice energy via polarization effects. For ionic compounds made of molecular cations and/or anions, there may also be ion-dipole and dipole-dipole interactions if either molecule has a molecular dipole moment. The theoretical treatments described below are focused on compounds made of atomic cations and anions, and neglect contributions to the internal energy of the lattice from thermalized lattice vibrations.

Born–Landé equation

Main article: Born–Landé equation

In 1918[5] Born and Landé proposed that the lattice energy could be derived from the electric potential of the ionic lattice and a repulsive potential energy term.[2]


NA is the Avogadro constant;
M is the Madelung constant, relating to the geometry of the crystal;
z+ is the charge number of the cation;
z is the charge number of the anion;
e is the elementary charge, equal to 1.6022×10−19 C;
ε0 is the permittivity of free space, equal to 8.854×10−12 C2 J−1 m−1;
r0 is the nearest-neighbor distance between ions; and
n is the Born exponent (a number between 5 and 12, determined experimentally by measuring the compressibility of the solid, or derived theoretically).[6]

The Born–Landé equation above shows that the lattice energy of a compound depends principally on two factors:

Barium oxide (BaO), for instance, which has the NaCl structure and therefore the same Madelung constant, has a bond radius of 275 picometers and a lattice energy of −3054 kJ/mol, while sodium chloride (NaCl) has a bond radius of 283 picometers and a lattice energy of −786 kJ/mol. The bond radii are similar but the charge numbers are not, with BaO having charge numbers of (+2,−2) and NaCl having (+1,−1); the Born–Landé equation predicts that the difference in charge numbers is the principal reason for the large difference in lattice energies.

Closely related to this widely used formula is the Kapustinskii equation, which can be used as a simpler way of estimating lattice energies where high precision is not required.[2]

Effect of polarization

For certain ionic compounds, the calculation of the lattice energy requires the explicit inclusion of polarization effects.[7] In these cases the polarization energy Epol associated with ions on polar lattice sites may be included in the Born–Haber cycle. As an example, one may consider the case of iron-pyrite FeS2. It has been shown that neglect of polarization led to a 15% difference between theory and experiment in the case of FeS2, whereas including it reduced the error to 2%.[8]

Representative lattice energies

The following table presents a list of lattice energies for some common compounds as well as their structure type.

Compound Experimental Lattice Energy[1] Structure type Comment
LiF −1030 kJ/mol NaCl difference vs. sodium chloride due to greater charge/radius for both cation and anion
NaCl −786 kJ/mol NaCl reference compound for NaCl lattice
NaBr −747 kJ/mol NaCl weaker lattice vs. NaCl
NaI −704 kJ/mol NaCl weaker lattice vs. NaBr, soluble in acetone
CsCl −657 kJ/mol CsCl reference compound for CsCl lattice
CsBr −632 kJ/mol CsCl trend vs CsCl like NaCl vs. NaBr
CsI −600 kJ/mol CsCl trend vs CsCl like NaCl vs. NaI
MgO −3795 kJ/mol NaCl M2+O2- materials have high lattice energies vs. M+O. MgO is insoluble in all solvents
CaO −3414 kJ/mol NaCl M2+O2- materials have high lattice energies vs. M+O. CaO is insoluble in all solvents
SrO −3217 kJ/mol NaCl M2+O2- materials have high lattice energies vs. M+O. SrO is insoluble in all solvents
MgF2 −2922 kJ/mol rutile contrast with Mg2+O2-
TiO2 −12150 kJ/mol rutile TiO2 (rutile) and some other M4+(O2-)2 compounds are refractory materials

See also



  1. ^ a b Atkins; et al. (2010). Shriver and Atkins' Inorganic Chemistry (Fifth ed.). New York: W. H. Freeman and Company. ISBN 978-1-4292-1820-7.
  2. ^ a b c David Arthur Johnson, Metals and Chemical Change, Open University, Royal Society of Chemistry, 2002,ISBN 0-85404-665-8
  3. ^ Zumdahl, Steven S. (1997). Chemistry (4th ed.). Boston: Houghton Mifflin. pp. 357–358. ISBN 978-0-669-41794-4.
  4. ^ Haynes, William M.; Lide, David R.; Bruno, Thomas J. (2017). CRC handbook of chemistry and physics : a ready-reference book of chemical and physical data. Boca Raton, FL: CRC Press, Taylor & Francis Group. pp. 12–22 to 12–34. ISBN 9781498754293.
  5. ^ I.D. Brown, The chemical Bond in Inorganic Chemistry, IUCr monographs in crystallography, Oxford University Press, 2002, ISBN 0-19-850870-0
  6. ^ Cotton, F. Albert; Wilkinson, Geoffrey; (1966). Advanced Inorganic Chemistry (2d Edn.) New York:Wiley-Interscience.
  7. ^ M. Birkholz (1995). "Crystal-field induced dipoles in heteropolar crystals I: Concept". Z. Phys. B. 96 (3): 325–332. Bibcode:1995ZPhyB..96..325B. CiteSeerX doi:10.1007/BF01313054. S2CID 122527743.
  8. ^ M. Birkholz (1992). "The crystal energy of pyrite". J. Phys.: Condens. Matter. 4 (29): 6227–6240. Bibcode:1992JPCM....4.6227B. doi:10.1088/0953-8984/4/29/007. S2CID 250815717.