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Ionic Bonding - Characteristics & Summary

Key Facts & Summary:

  • Ionic bond, is the force of attraction between oppositely charged species (ions).
  • An ion is an atom that has a non-zero net electrical charge.
  • Whether an element is going to be the cation or anion in an ionic bond depends on several factors, the periodic table can serve as a guide.
  • Elements at the right of the periodic table tend to gain electrons to reach the electron configuration of the next higher noble gas.
  • Solid materials with ionic bonds are hard, good insulators, transparent, rigid and with high melting point.

When atoms combine to each other they create a chemical bond mainly driven by the attractive force between atoms. One type of chemical bond, called an ionic bond, is the force of attraction between oppositely charged species (ions). An ion is an atom that has a non-zero net electrical charge.

By convention the charge of an electron is considered negative and it is opposite to the proton charge that is considered positive.

Since the charge of the electron (considered "negative" by convention) is equal and opposite to that of the proton (considered "positive" by convention), the net charge of an ion is non-zero due to its total number of electrons being unequal to its total number of protons. A cation is a positively charged ion, with fewer electrons than protons, while an anion is negatively charged, with more electrons than protons. Because of their opposite charges, cations and anions attract each other and readily form ionic compounds.

Whether an element is the source of the cation or anion in an ionic bond depends on several factors, for which the periodic table can serve as a guide.

An atom with a complete shell of valence electrons (corresponding to 8 electrons) tends to be chemically inert. Atoms with one or two less valence electrons are highly reactive because a low amount of energy is needed to remove the extra electron. It requires relatively low energy to remove the extra valence electrons to form a positive ion. Furthermore, they have tendency to gain electrons to have a complete shell (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond).

Similar to an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as atomic excitation. Or the electron can even break free from its associated atom's valence shell; this is ionization to form a positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which is not fully occupied.

Valence energy levels correspond to the principal quantum numbers (n = 1, 2, 3, 4, 5 ...) or are labeled alphabetically with letters used in the X-ray notation (K, L, M, …).

In forming ionic compounds, elements at the left of the periodic table typically lose electrons, forming a cation that has the same electron configuration as the nearest noble gas. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a metal, and an electron acceptor, a nonmetal. Loss of an electron from sodium, for example, gives the species Na, which has the same electron configuration as neon.

Ionic bonding is the complete transfer of valence electron(s) between atoms. The electrons in the outermost shell are the valence electrons, the electrons on an atom that can be gained or lost in a chemical reaction. Ionic bonding is observed because metals have few electrons in their outer-most orbitals. By losing those electrons, these metals can achieve noble gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonding, more than 1 electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the number of electrons donated or received. In ionic bonds, the net charge of the compound must be zero.

In this example, the magnesium atom is donating both of its valence electrons to chlorine atoms.

Each chlorine atom can only accept 1 electron before it can achieve its noble gas configuration; therefore, 2 atoms of chlorine are required to accept the 2 electrons donated by the magnesium. Notice that the net charge of the compound is 0.

Characteristics

There are several characteristic that belongs to most of the compounds that are solids and their atoms are keep together by ionic bonds.

Generally, solid materials with ionic bonds are hard, they resist to localized plastic deformation. This is due to the fact that particles cannot easily slide past one another.

These materials are good insulators. It means that internal electric charges do not flow freely and very little electric current will flow through it under the influence of an electric field. This effect is happening because there are no free electrons.

Solid ionic materials are usually transparent because their electrons are not moving from atom to atom and they do not interact with light photons.

They often are having hardness and rigidity but little tensile strength. They can break readily with a comparatively smooth fracture, as glass and they do not deform.

Solid ionic materials generally have high melting point because ionic bonds are relatively strong.

Here are examples of ionic bonds and ionic compounds:

NaBr - sodium bromide

KBr - potassium bromide

NaCl - sodium chloride

NaF - sodium fluoride

KI - potassium iodide

KCl - potassium chloride

CaCl2 - calcium chloride

 

References and further readings:

https://www.slideserve.com/

https://chem.libretexts.org/Textbook_Maps/Organic_Chemistry/Supplemental_Modules_(Organic_Chemistry)/Fundamentals/Ionic_and_Covalent_Bonds

https://www.youtube.com/watch?v=Qf07-8Jhhpc

http://www.chemistryrocks.net

“Organic chemistry”, Francis A. Carey, ISBN 0-07-117499-0