Ionic Bond Definition and Examples (2024)

An ionic bond or electrovalent bond is an electrostatic attraction where one atom donates an electron to another atom. The transfer results in the atom that loses an electron become a positively charged ion or cation, while the atom gaining the electron becomes a negatively charged ion or anion. But, the net charge on an ionic compound is zero (neutral). This type of chemical bond occurs between atoms with very differently electronegativity values, such as metals and nonmetals or various molecular ions. Ionic bonding is one of the main types of chemical bonding, together with covalent bonding and metallic bonding.

• An ionic bond is when one atom donates its valence electron to another atom, increasing the stability of both atoms.
• This type of bond forms when atoms or molecular ions have electronegativity differences greater than 1.7.
• Ionic bonds produce compound that conduct electricity when dissolved or molten and generally have high melting and boiling points as solids.
• Because of the polarity of the chemical bond, many ionic compounds dissolve in water.

Examples of Ionic Bonds

The classic example of an ionic bond is the chemical bond that forms between sodium and chlorine atoms, forming sodium chloride (NaCl). Sodium has one valence electron, while chlorine has seven valence electrons. When a sodium atom donates its lone electron to chlorine, the sodium gains a +1 charge, but becomes more stable because its electron shells are complete. Similarly, when chlorine accepts an electron from sodium, it gets a -1 charge and completes the octet of its valence electron shell. The resulting ionic bond is very strong because there is no repulsion between neighboring electrons, like you see when atoms share electrons in a covalent bond. That being said, covalent bonds can also be strong, as when carbon atoms share four electrons and form diamond.

Another example of an ionic bond occurs between magnesium and hydroxide ions in magnesium hydroxide (MgOH₂). In this case, the magnesium ion has two valence electrons in its outer shell. Meanwhile, each hydroxide ion gains stability if it gains an electron. So, magnesium donates one electron to one hydroxide and one electron to the other hydroxide, giving the Mg atom a +2 charge. The hydroxide ions then each have a charge of -1. But, the compound is neutral. You only see Mg²⁺ and OH^– in solution or when the compound is molten. Note the chemical bond between the oxygen and hydrogen in hydroxide is covalent.

Here are other examples of compounds containing ionic bonds:

• Potassium chloride, KCl
• Magnesium sulfate, MgSO₄
• Lithium chloride, LiCl
• Cesium fluoride, CeF
• Strontium hydroxide, Sr(OH)₂
• Potassium Cyanide, KCN

Properties of Ionic Compounds

Compounds that contain ionic bonds share some common properties:

• They are usually solid at room temperature.
• Ionic compounds are electrolytes. That is, they conduct electricity when dissolved or molten.
• They typically have high melting and boiling points.
• Many ionic compounds are soluble in water and insoluble in organic solvents.

Predicting an Ionic Bond Using Electronegativity

Atoms or ions with large electronegativity differences form ionic bonds. Those with small or no electronegativity differences form covalent bonds, unless they are metals, in which case they form metallic bonds. The values for the electronegativity differences vary according to different sources, but here are some guidelines for predicting bond formation:

• An electronegativity difference greater than 1.7 (1.5 or 2.0 in some texts) leads to ionic bonding.
• A difference greater than 0.5 (0.2 in some texts) and less than 1.7 (or 1.5 or 2.0) leads to polar covalent bond formation.
• An electronegativity difference of 0.0 to 0.5 (or 0.2, depending on the source) leads to nonpolar covalent bond formation.
• Metals bond to each other via metallic bonding.

But, in all of these bonds there is some covalent character or sharing of electrons. In an ionic compound, for example, there is no “clean” ionic bonding or total transfer of electrons (even though it is drawn that way in diagrams). It is just that the bond is much more polar than in a covalent bond. Similarly, in metallic bonding, some association exists between a metallic nucleus and the mobile valence electrons.

Also, be aware there are many exceptions to these guidelines. Many times the electronegativity difference between a metal and nonmetal is around 1.5, yet the bond is ionic. Meanwhile, the electronegativity difference between hydrogen and oxygen (a polar covalent bond) is 1.9! Always consider whether the participating atoms are metals or nonmetals.

Example Problems

(1) What type of chemical bond forms between iron (Fe) and oxygen (O)?

An ionic bond forms between these two elements. First, iron is a metal and oxygen is a nonmetal. Second, their electronegativity values are significant (1.83 for iron and 3.44 for oxygen).

(2) Which of these two compounds contains ionic bonds? CH₄ or BeCl₂

BeCl₂ is the ionic compound. CH₄ is a covalent compound. The quick way of answering the question is looking at the periodic table and identifying which atoms are metals (Be) and which are nonmetals (H, Cl). A metal bonding to a nonmetal forms an ionic bond, while two nonmetals form a covalent bond. Otherwise, consult a chart of electronegativity values. The difference between the electronegativities of C and H is small, while the difference between Be (1.57) and Cl (3.16) is large (1.59). (Note this electronegativity difference, by itself, might lead you to predict a polar covalent bond. So, always look at whether atoms are metals or nonmetals.)

References

• Atkins, Peter; Loretta Jones (1997).Chemistry: Molecules, Matter and Change. New York: W.H. Freeman & Co. ISBN 978-0-7167-3107-8.
• Lewis, Gilbert N. (1916). “The Atom and the Molecule”.Journal of the American Chemical Society. 38 (4): 772. doi:10.1021/ja02261a002
• Pauling, Linus (1960). The Nature of the Chemical Bond and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry. ISBN 0-801-40333-2. doi:10.1021/ja01355a027
• Wright, Wendelin J. (2016). The Science and Engineering of Materials (7th ed.). Global Engineering. ISBN 978-1-305-07676-1.

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