# Structural Biochemistry/Chemical Bonding/Ionic interaction

An Ionic Bond is a specific type of chemical bond formed between a "metal" and a "nonmetal." "Metals" involved in ionic bonds are usually the Alkali and Alkaline-Earth metals - also known as the first two columns on the period table - as well as several transition state metals. The non-metals usually involved in the ionic bonds are the halogens.
The goal of forming chemical bonds is to have an octet. "Octet" means that an element has eight electrons. Another way of putting this phenomenon is to state that each element wants to have the electron configuration of a noble gas. For the non-metals, the goal is to achieve the electron configuration of the noble gas in the same row. For the metals, the goal is to have the electron configuration of the noble gas in the row directly preceding it. For example, chlorine wants the electron configuration of Argon while sodium wants that of Neon. Ionic bonds help achieve the octet because the metal effectively transfers its valence electrons to the nonmetal. In this way, both metal and nonmetal achieve noble-gas electron configurations.
However, pure ionic bonds do not exist. There is a level of covalent-bond character in each ionic bond. As a general rule of thumb, the larger the electronegativity difference between the metal and nonmetal, the more ionic the bond, and therefore, the less covalent the bond. Electronegativity is the ability to draw electron density while in the presence of another atom. The more electronegative the atom, the greater its ability to pull electron density towards itself.
Electrostatic interactions are also known as Charge-charge interactions and Ionic interactions. An electrostatic attraction exists when there are closely packed ions of opposite charges. An electrostatic repulsion is present between different ions that have the same charge.

## Coulomb's Law

The force between two point charges can be calculated by Coulomb’s law,
F$_{1on2}$ = F$_{2on1}$ = kq$_1$q$_2$/r$^2$
In other words, the bond energy is directly proportional to the charges of the two atoms and inversely proportional to the square of the distance between the two atoms. F has the unit N (Newtons), r is the distance between the center of the two-point charges in meter, q$_1$ and q$_2$ are the charges (in C, Coulombs) of each atom respectively. k is the constant, approximately equal $9.00x10^{9}$ N•m$^2$•C$^{−2}$. If the force is negative (F < 0), it represents the existence of attraction since the only possible way to have F smaller than zero is if the sign of q$_1$ and q$_2$ are different, or opposite, meaning the sign of charges for these two atoms are opposite. In other words, if the force is positive (F > 0), it means the two charged atoms repel each other, due to the same sign of charges (both positive or negative charges.)

## Ionic bond energy

The ionic bond energy between two-points charges can be calculated using the formula of electrical potential energy,
U$_{ele}$= Kq$_1$q$_2$/r$^2$
In other words, the bond energy is directly proportional to the charges of the two atoms and inversely proportional to the distance (bond length in microscopic level) between the two charged atoms. U$_{ele}$ has the unit J (Joules), r is the distance between the ion centers in nanometers, and q$_1$ and q$_2$ are the numerical ion charges. K is the constant, approximately equal $2.31x10$$^{−19}$ J•nm. It is also equal to kQ$^2$ x 10$^{−9}$ (Conversion to nanomemter, nm) or Q$^2$/4$\pi \epsilon _{0}$ x 10$^{−9}$ where k is approximately equal to $9.00x10^{9}$ Nm$^2$C$^{−2}$ and Q is equal to the charge of a proton, $1.60x10$$^{−19}$C; $\epsilon _{0}$ approximately equal to $8.85x10$$^{−12}$ C$^2$N$^{−1}$•m$^{−2}$ or F•m$^{−1}$, representing the permittivity of free space.

## Lattice Energy

Lattice energy (ΔH$^0$$_{lattice}$) is the released energy from the interaction between ions. The positive sign of this energy indicates that the energy is needed for ions to form a solid. On the other hand, the negative sign of the energy indicates that the energy is needed for ionic solids to be separated into its gaseous ions. In addition, ionic interactions can be explained based on the lattice energy because the qualitative number of the energy indicates the ions' hardness, solubility, and melting point. Lattice energy can be predicted based on the effect of ionic size and ionic change. As one goes down the periodic table, the ionic radius increases. And as the radius increases, there would be a decrease in the electrostatic energy between the positively charged atom and negatively charged atom. As a result, there would be a decrease in the lattice energy. In addition, ionic charge can also determine the lattice energy. A greater amount of energy will be needed to bring a larger charged ions together than to form the smaller charged ions.

## References

* ↑ Silberberg, Martin S.(2010). Principles of General Chemistry (2nd Edition).McGraw Hill Publishing Company. ISBN 978-0-07-351108-05