Ionic Bond-Definition, Types, Properties & Examples-Embibe
Create free account
  • Written By Sushmita Rout
  • Last Modified 24-06-2022
  • Written By Sushmita Rout
  • Last Modified 24-06-2022

Ionic Bond: Definition, Types, Properties, Examples

The ionic bond is the electrostatic attraction between oppositely charged ions in a chemical molecule. Iconic bonds are also known as electrovalent bonds.  What substances are used to make lemonade, lemon juice, sugar, water, and obviously salt? Do you know how salt is different from sugar? Why is oil not miscible in water? Why does salt appear as crystals but not water?

All this happens because the atoms of each compound are differently bonded to each other. In this article, we shall discuss the ionic bond definition, covalent compounds, ionic bond examples and types of bond chemistry.

What is Ionic Bond?

According to the octet rule, an atom is most stable when there are eight electrons in its valence shell. To a solid-state attain stability, atoms lose, gain or share electrons present in their valence shell. An atom that loses one or more valence electrons to become a positively charged ion forms a cation, while an atom that gains electrons and becomes a negatively charged ion forms an anion.

When there is a complete transfer of electrons between the positively charged cation and the negatively charged anion, an electrostatic force of attraction develops, known as the Ionic Bond. This exchange of valence electrons allows ions to achieve electronic configurations of the neighbouring noble gases, satisfying the octet rule. A cation is represented by a positive superscript charge \(\left(  +  \right)\) to the right of the atom.

Ionic Bond

Learn Exam Concepts on Embibe

An anion is represented by a negative superscript charge \(\left(  –  \right)\) to the right of the atom.

Formation of Anion

For example, when a sodium atom loses one electron, it will have one more proton more than an electron, rendering it an overall positive \(\left( { + 1} \right)\) charge. For the sodium ion, the chemical symbol is \({\rm{N}}{{\rm{a}}^{ + 1}}\) or just \({\rm{N}}{{\rm{a}}^ + }.\)

Similarly, a chloride ion \({\rm{C}}{{\rm{l}}^ – }\) is formed if a chlorine atom gains an extra electron. These ionic species are more stable than the atom due to the octet rule.

Forming an Ionic Bond

The chemical bond that is formed between \(2\) atoms through the transfer of one or more electrons from the electropositive or metallic element to the atom of an electronegative or non-metallic element is called an ionic or electrovalent bond.

Ionic Bond Formation

Practice Exam Questions

We know that the electronic configuration of the sodium atom is \(2,8,1\). It has only one electron in its outermost shell. By donating this electron, it acquires the inert gas electronic configuration of Neon \(\left( {2,8} \right).\)

electronic configuration of Neon

On the other hand, the electronic configuration of the chlorine atom is \(2,8,7\). It needs only one more electron to complete its octet to acquire the inert configuration of \(\left( {2,8,8} \right)\) of Argon.

inert configuration of Argon

Here, the sodium atom requires an adequate amount of energy equal to its ionization energy of about \(496\,{\rm{kJ}}/{\rm{mol}}\) to remove an electron from its outermost shell to form a positively charged sodium ion \({\rm{N}}{{\rm{a}}^ + }.\) As energy is consumed in this process, it is called an endergonic process.

On the contrary, the chlorine atom, deficient of one electron, accepts an electron; it releases energy equal to its electron affinity of the chlorine atom,i.e. \(349\,{\rm{kJ}}/{\rm{mol}},\) to give a chloride ion. Since energy is released in this process, it is known as the Exorgenic process.

Exorgenic Process

Thus, the oppositely charged ions formed have strong forces of attraction called the electrostatic forces of attraction. These forces bring the ion closer to form an ionic bond. Thus, the electrostatic forces form the basis of an ionic bond.

Electrovalency– The number of atoms lost by one atom or gained by the other atom is called electrovalency.

Electrovalency of sodium and chlorine in \({\rm{NaCl}}\)  is one. Hence, they are monovalent.

Similarly, in the formation of Calcium oxide \(\left( {{\rm{CaO}}} \right),\) calcium donates two valence electrons to form calcium ions \(\left( {2,8,8} \right),\) and oxygen gains two electrons to form oxide ions \(\left( {2,8} \right).\) Thus the electrovalency of calcium and oxygen in \(2\) each, i.e. they are divalent.

Once the transfer of electrons forms the oppositely charged ions. The ionic bond formation between sodium and fluorine atoms is shown below.

Ionic bond formation of sodium and flourine

Attempt Mock Tests

Writing Formula of an Ionic Compound

To determine the chemical formulas of ionic compounds, the following two conditions must be fulfilled:

  1. The cation and the anion should follow the octet rule for maximum stability.
  2. Ions should combine in a way that the charges of the ions must balance out and the overall ionic compound is neutral.

The charges present 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.

Ionic Compound Formula

Conditions for Formation of Ionic Bond

In ionic compounds, atoms need to fulfil several conditions to form an ionic bond. We are all quite familiar with table salt that we use to add flavour to food. We also know that chemically, table salt is sodium chloride in which sodium and chloride ions are bonded together by ionic bonds. Various factors affect the formation of these ions and consequently ionic bonding, which ultimately gives rise to ionic compounds. The factors are –

(a) The number of valence electrons present in the atoms involved in bonding -.
Atoms possessing one, two, or three valence electrons which are positioned in groups \({\rm{I}}\left( {\rm{A}} \right){\rm{,II}}\left( {\rm{A}} \right){\rm{,}}\) and \(13\left( {\rm{A}} \right)\) of the modern periodic table, tend to lose electrons and form positively charged species called cations.

On the other hand, atoms having \(5,{\rm{ }}6,\) or \(7\) valence electrons present in groups \(15,{\rm{ }}16,\) and \(17\) of the modern periodic table, have a greater tendency to accept electrons and form negatively charged species called anions.

Thus, an element like potassium which belongs to group \({\rm{I}}\left( {\rm{A}} \right)\) of the modern periodic table and has one electron in its outermost shell, is ideal for forming an ionic bond with Chlorine which belongs to group \(17\) and has \(7\) electrons in its outermost shell. Potassium Chloride thus, form is, therefore, an ionic compound.

Potassium Chloride

(b) The low ionization energy of the metal – The ionisation energy is the minimum energy required to remove an electron from the outermost shell of a neutral gaseous atom.

Consider the formation of sodium ion \(\left( {{\rm{N}}{{\rm{a}}^ + }} \right)\)  from sodium atom. The ionization energy of sodium is about \(500\,{\rm{kJ}}/{\rm{mol}},\) which is quite low. It can easily lose electrons and get converted into a sodium ion. This ion can further take part in ionic bond formation with other anions such as \({\rm{C}}{{\rm{l}}^ – },{\rm{B}}{{\rm{r}}^ – }.\) Thus, the lower value of ionization energy of metal will favor the ionic bond formation.

(c) The electron affinity of a non-metal – By definition, electron affinity is the energy released when an electron is added to a neutral isolated gaseous atom.

Consider an atom of fluorine; it has \(7\) electrons in its outermost orbit. It readily accepts one electron to complete its octet configuration. This process releases about \(328\,{\rm{kJ}}\) of energy per mol. The fluoride ion thus formed possesses lower energy than the fluorine atom.

We know that lower energy corresponds to greater stability. Hence, the fluoride ion is more stable than the fluorine atom. Thus, the higher electron affinity of a non-metal will favour the formation of an anion and thereby will lead to a  stable ionic compound.

(d) The lattice energy of the ionic compound – An important factor that affects the stability of an ionic compound is the lattice energy of the ionic compounds. Lattice energy is the energy released when one gram mole of a crystal is formed from its gaseous ions. In any crystal, the constituent ions of the ionic compound are held together by electrostatic forces of attraction.

lattice energy of the ionic compound

The stronger the forces of attraction, the higher is the lattice energy and the more stable is the compound. This electrostatic force of attraction is determined by Coulomb’s Law.

Coulomb’s law states that the force of attraction \(\left( {\bf{F}} \right)\) between two oppositely charged ions (\({\bf{Q1}}\) and \({\bf{Q2}}\)), which are separated in air by the distance \({\rm{d,}}\) is given by the following formula which is as follows-

\({\rm{F = K}}\frac{{{{\rm{Q}}_{\rm{1}}}{{\rm{Q}}_{\rm{2}}}}}{{{{\rm{d}}^{\rm{2}}}}}\)

Where d is the distance of the radii of the cation \(\left( {{{\rm{r}}^{1 + }}} \right)\) and radii of the anion \(\left( {{{\rm{r}}^{2 – }}} \right),\) \({\bf{K}}\) is the constant of proportionality. This relationship indicates that the attractive force depends on the magnitude of the charges (\(Q1\)and \(Q2\)) and the distance between these two charges. Thus, we can say that that the magnitude of lattice energy depends on the following two factors.

  1. The magnitude of the charges of the ions.
  2. Size of the ions

(i) Let us consider two ionic compounds, Magnesium oxide, and sodium chloride. It is assumed that the distance between these compounds is the same.

Magnesium oxide is a divalent compound, i.e. both the magnesium cation and oxygen anion have a charge of \(2\). Whereas sodium chloride being a monovalent compound, both the sodium cation and chloride anion carry a charge of \(1\). Applying coulomb’s law for both the compounds we get-

\({\rm{NaCl}};{\rm{F}} = \frac{{{\rm{k}}(1 \times 1)}}{{{{\rm{d}}^2}}} = {\rm{k}}\left( {\frac{1}{{{{\rm{d}}^2}}}} \right)\)

\({\rm{MgO}};{\rm{F}} = \frac{{{\rm{k}}(2 \times 2)}}{{{{\rm{d}}^2}}} = {\rm{k}}\left( {\frac{4}{{{{\rm{d}}^2}}}} \right)\)

This means that the force of attraction between \({\rm{M}}{{\rm{g}}^{2 + }}\) ions and \({{\rm{O}}^{2 – }}\) ions in a crystal of magnesium oxide is \(4\) times greater than that unipositive sodium and uni negative chloride ion in a crystal of sodium chloride. This explains the higher stability of \({\rm{MgO}}\) compared to \({\rm{NaCl}}{\rm{.}}\) Thus, the Higher the magnitude of the charges of ions, the higher is its lattice energy, the more stable is the ionic compound

(ii) Let us consider \({\rm{NaCl}}\) and \({\rm{CsCl}}.\) These two ionic compounds are univalent, i.e. the charges of the cations and anions are one. But, the ionic radius of \({\rm{Na}}\) is smaller as compared to the \({\rm{Cs}}\) ion. Due to the smaller radius of the sodium ion, both sodium and chloride ions are closer to each other compared to \({\rm{CsCl}}.\) Thus, the force of attraction between the ions in \({\rm{NaCl}}\) is greater than \({\rm{CsCl}}.\) This makes sodium chloride more stable than Cesium chloride.

(e) The difference in electronegativity between two atoms – The electronegativity of an atom is a measure of an atom’s ability to attract electrons towards itself involved in a bond formation.

 If the difference in electronegativity values of the atoms is \(1.7\) or more, then the formation of an ionic bond is favoured.

Consider the molecules of \({\rm{NaCl}}\) and  \({\rm{HCl}}{\rm{.}}\)

For \({\rm{NaCl;}}\) Electronegativity difference \(= 3.0\left( {{\rm{Cl}}} \right) – \;0.9\left( {{\rm{Na}}} \right) = 2.1.\)

This helps the formation of a stable ionic bond in sodium chloride.

For \({\rm{HCl;}}\) Electronegativity difference \( = 3.0\left( {{\rm{Cl}}} \right) – \;2.1\left( {\rm{H}} \right) = 0.9\)

Due to the smaller electronegativity difference, the bond between hydrogen and chlorine is covalent.

Properties of Ionic Bond

Compounds composed of cations and anions are called ionic compounds:

CompoundCationAnion
Sodium Chloride \(\left( {{\rm{NaCl}}} \right)\)\({\rm{N}}{{\rm{a}}^ + }\)\({\rm{C}}{{\rm{l}}^ + }\)
Calcium oxide \(\left( {{\rm{CaO}}} \right)\)\({\rm{C}}{{\rm{a}}^{2 + }}\)\({{\rm{O}}^{2 – }}\)
Magnesium oxide \(\left( {{\rm{MgO}}} \right)\)\({\rm{M}}{{\rm{g}}^{2 + }}\)\({{\rm{O}}^{2 – }}\)

In all these compounds, oppositely charged ions are held together by a strong electrostatic force of attraction. When an electropositive element like \({\rm{N}}{{\rm{a}}^ + }\) loses its electron present in the valence shell to an electronegative element \(\left({{\rm{C}}{{\rm{l}}^ – }} \right),\) an ionic bond is formed.

Both the atoms thus achieve a stable octet configuration. Similarly, metal-like calcium loses two of its outermost electrons to a non-metal-like solid-state, oxygen resulting in an ionic bond between calcium and oxygen forming Calcium oxide \(\left({{\rm{CaO}}} \right).\)

Calcium oxide

1. A metal always forms the cation, whereas a non-metal always forms the anion

cation and anion

2. Most of the ionic compounds are crystalline solids at room temperature – The constituent ions of the ionic compound attract one another strongly and are arranged in a repeating three-dimensional pattern. It is this arrangement that gives the crystal a characteristic geometrical shape.

Ionic bond formation of metal and nonmetal

3. Ionic compounds generally have high melting points – The ions present in ionic compounds are held together by very strong attractive forces. Ionic compounds have very high melting points, i.e. they need a lot of heat energy to break the bond between them.

4. Electrical conductivity – In solid-state, ionic compounds are generally non-conductors of electricity. When heated to a temperature above their melting point, the electrostatic force of attraction between the ions breaks, and the ions become free to move. These free ions can now allow the passage of electricity.

5. Solubility in water – When an ionic compound like sodium chloride is added to water, the negative end of the water molecule attracts the cations and pulls them out of the crystal. Likewise, The positive end of the water molecule pulls the anions resulting in the dissolution of the compound in water. Thus, ionic compounds are soluble in polar solvents like water.

Solubility of Ionic Compound

6. Ionic compounds are brittle – When an external force is applied to the crystals of an ionic compound, it shatters into pieces. This happens because, in the crystals of sodium chloride, the \({\rm{N}}{{\rm{a}}^ + }\) ions and \({\rm{C}}{{\rm{l}}^ – }\) ions are lined up against each other in a lattice with a strong electrostatic force of attraction.

When an external force is applied, the alignment of the ions changes in a way that the like charges come close together. This results in a strong electrostatic repulsion, and the ions move apart. As the ions shatter into pieces, the shape of the crystals breaks.

Ionic Bond vs Covalent Bond

Ionic Bond vs Covalent Bond
ParametersCovalent BondIonic Bond
What is it?It is a form of chemical bonding between two non-metallic atoms, which is characterized by the sharing of pairs of electrons between atoms and other covalent bonds.Also known as an electrovalent bond, it is a type of bond formed from the strong electrostatic force of attraction between oppositely charged ions in a chemical compound.
Occurs betweenTwo non-metals or a non-metal and a metalloidOne metal and one non-metal
Nature of component particlesIt consists of electrically neutral discrete moleculesIonic compounds are made up of oppositely charged particles called cations and anions.
State at room temperatureGases, liquids, or low melting solidsCrystalline solids
PolarityLowHigh
SolubilityPolar covalent compounds dissolve in polar solvents. Ex- HCl in water. Non-polar covalent compounds are soluble in organic liquids only.Being polar, ionic compounds are soluble in polar solvents only like water. Not soluble in organic liquids
FormationThey are formed between two non-metal having similar electronegativities. Neither of the atoms is strong enough to attract electrons from the other. For stabilization, neighbouring they share their electrons from the outermost orbital.They are formed between a metal(+ion) and a non-metal (-ve ion). Non-metals are stronger than metals and can get electrons very easily from metals. These two opposite ions attract each other and form an ionic bond.
ConductivityNon-polar covalent compounds do not dissociate into ions. Hence, do not conduct electricity. Polar covalent compounds easily dissociate in water and behave as good conductors of electricity. Ex- HCl in water.Ionic compounds conduct electricity when molten or dissolved in water.
Melting PointLowHigh
Boiling PointLowHigh
ExamplesMethane, Hydrochloric acidSodium Chloride, Sulphuric acid
Ionic Bond and Covalent Bond

Covalent Character of an Ionic Bond – Fajan’s Rule

Carbon dioxide, water, and chlorine gas are some common examples of compounds having a covalent bond. On the other hand, few compounds like table salt, magnesium oxide, and calcium chloride are ionic. But in reality, no bond or compound is completely ionic or covalent in nature. In an ionic bond, oppositely charged ions are formed when the electropositive element donates valence electrons to the electronegative element.

When these ions approach each other, due to the increased nuclear charge on the cation, it attracts the loosely bound outermost electrons of the large anion. As a result, the electron cloud of the anion gets distorted by its neighbouring cation. This process is known as Polarisation, and the ability of the cation to attract the electron cloud of the anion is called its polarising power. The tendency of the anion to get polarised by the cation is called polarizability. 

Charge and size of the cation

For example – In the compounds, \({\rm{NaCl}}\) and \({\rm{CaC}}{{\rm{l}}_2},\) the ionic radii of sodium and calcium are very close. However, due to the presence of a higher nuclear charge, the \({\rm{C}}{{\rm{a}}^{2 + }}\) ion can polarise the electron cloud of chloride much more strongly than \({\rm{N}}{{\rm{a}}^ + }.\) Thus, \({\rm{CaC}}{{\rm{l}}_2},\) is more covalent than \({\rm{NaCl}}.\)

When \({\rm{CaC}}{{\rm{l}}_2}\) is compared with \({\rm{BeC}}{{\rm{l}}_2},\) it was observed that the cation \({\rm{B}}{{\rm{e}}^{2 + }}\) is much smaller in size than \({\rm{C}}{{\rm{a}}^{2 + }}\) can polarise the electron cloud of chloride to a great extent. This happens because the positive charge on the beryllium ion is concentrated in a small area, whereas the same charge is distributed over a large surface area in calcium ions. Hence, beryllium chloride is more covalent than calcium chloride.

Charge and size of cation

Size of the Anion

For example – In the compounds LiF and LiI, the iodide ion is larger as compared to the fluoride ion. The outermost electrons are well shielded and held more loosely by the nucleus, making it easily polarisable by the cations. This is not the case with fluoride ions. As its size is smaller than the iodide ion, its electron is held strongly by its nucleus. Hence, LiI has a more covalent character than LiF.

Fajan's Rules

Cations with pseudo noble gas configuration (\(18\) electrons in the outermost shell), as in the case of \({\rm{CuC}}{{\rm{l}}_2},\) have higher polarising power than cations having true, noble gas configuration \(\left( {{\rm{N}}{{\rm{a}}^ + }} \right).\) Electrons in the \({\rm{C}}{{\rm{u}}^{2 + }}\) are not effectively shielded by its nucleus.

As a consequence, the cation exerts a higher polarising power over chloride ions and polarises its electron cloud to a greater extent. Although the cations in both the compounds have similar charge and radius, \({\rm{CuC}}{{\rm{l}}_2}\) have more covalent character than \({\rm{NaCl}}.\)

Examples of Compounds Having Ionic Bond

Ionic Bond Examples

Summary

Compounds containing ionic bonds are part and parcel of our day-to-day life, from the flavouring agent table salt to the tubes of toothpaste. From baking soda to washing soda, from bleaches to preservatives, all consist of compounds containing ionic bonds. Hence, to summarise, it is indeed needless to say that without ionic compounds, our survival is impossible.

FAQs on Ionic Bond

Q.1: Frame the formula for the compound formed when magnesium and fluorine combine?
Ans: Mg forms a divalent ion of \({\rm{M}}{{\rm{g}}^{2 + }}.\) To obey the octet rule, Mg gets rid of the two valence electrons present in its valence shell. Fluorine has seven valence electrons and would gain an electron to complete its octet configuration, forming an F- ion. As \({\rm{M}}{{\rm{g}}^{2 + }}\) needs two electrons to neutralize its charge, it combines with \(2\,{{\rm{F}}^ – }\) ions forming an ionic compound. Hence, the formula of the ionic compound formed is \({\rm{Mg}}{{\rm{F}}_2}.\) The subscript two indicates that two fluorines are ionically bonded to magnesium.

Q.2: Which elements most often form ionic bonds?
Ans: Ionic bonds most commonly form between metals and non-metals. There is a large electronegativity difference between metals and non-metals. If the difference is \(1.7\) or more, an ionic bond is favoured.

Q.3: Does oxygen form ionic bonds?
Ans: Oxygen does not contain ionic bonds. There is a mutual sharing of electrons between two oxygen atoms which results in the formation of a covalent bond. Sharing of electrons takes place to attain the noble gas configuration of argon \(\left( {2,8} \right).\)

Q.4: Is water covalent or ionic?
Ans: Water is a polar covalent molecule, which easily dissolves most of the substances. As oxygen is more electronegative than the hydrogen atom, the shared pair of electrons is more attracted towards the oxygen atom. This results in the formation of a partially positive charge over the hydrogen atom and a partial negative charge over the oxygen atom.

We hope that you might have now understood a lot about the Ionic Bond concept. If there is something that you get stuck on, feel free to comment below. And we will get back to you at the earliest.

Master Exam Concepts with 3D Videos