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  • Writer's pictureMatthew To

Chemistry 101: Ionic and Covalent Bonding

In chemistry, atoms often need to bond together to form more complex molecules and compounds to become “stable.”

In other words, when an atom is “stable,” it has 8 outermost electrons, otherwise called valence electrons. This can be achieved in two ways: Ionic Bonding and Covalent Bonding.


But first, what are valence electrons?

All bonding starts with the concept of valence electrons, which are electrons located on the outermost shell of the atom.

On the periodic table, an element’s position on the periodic table indicates the number of valence electrons present in its atom.

image of a periodic table with group numbers

An element like Lithium (Li) in Group 1 would have 1 valence electron, while an element such as Chlorine (Cl) in Group 17 would have 7 valence electrons.

Groups = Columns

We generally skip transition metals in Groups 3-17 because their outer shells and electron configurations are different and these metals often bond through another process called metallic bonding. Elements in Group 18 are called noble gases and tend to have little reactivity because they are already stable with 8 valence electrons.


Ionic Bonding:

Ionic bonding starts with ions, atoms that have either accepted or donated electrons based on the number of electrons in their valence shell.

For example, as we mentioned before, Sodium (Na) has 1 valence electron in its outermost shell. If possible, a Sodium atom would like to lose or donate this electron, so that it can more stably exist with a full valence shell from an inner layer.

Traditionally, atoms are neutral, but since we altered the amount of electrons and charge within the atom, we now can refer to it as an ion. Since it lost an electron, it now has a positive charge because there are more protons than electrons present in its atom. Positively charged ions are referred to as cations, and we would denote this positively charged ion with a +.

Now, let’s utilize the element Chlorine. It is in the 17th group in the periodic table and has 7 valence electrons. To reach a desirable full outer shell, Chlorine would only need to accept one valence electron to have 8. If we take the electron lost by Sodium and donate it to Chlorine, the Chlorine would now also obtain a charge. However, this time, the charge would be negative because it received an extra electron. We refer to negatively charged ions as anions and Chlorine would be written as Cl-.

In the ionic bonding process, an atom donates an electron and becomes a positively charged cation, and an atom accepts this electron to become a negatively charged anion. The two opposite charges attract to form a bond. Utilizing the two ions that we created above, the Na+ ion and the Cl- ion would bond together to form NaCl, a compound that we commonly refer to as table salt.

Remember, atoms always try to achieve stability when bonding with another atom. Hence, they either gain or lose electrons depending on which is more favorable. For example, Calcium in Group 2 of the periodic table would form a 2+ ion because losing 2 electrons will give it a full outer shell. Calcium can either bond with other anions with a 2- charge or with two 1- ions like Cl-.


Covalent Bonding

Covalent bonding occurs between nonmetals and is the process when two or more atoms share one or more pairs of electrons to form a compound.

Let's look at an example of covalent bonding between hydrogen (H) and oxygen (O). Hydrogen has 1 valence electron and oxygen has 6 valence electrons. Oxygen wants to gain 2 electrons to complete its outermost shell, while hydrogen wants to gain 1 electron to complete its outermost shell.

However, these atoms cannot ionically bond because they are both nonmetals as they cannot donate electrons to one another. Therefore, these atoms covalently bond, or share electrons instead. In this process, atoms share the electrons in their outermost valence shell.

When hydrogen and oxygen come into contact, they will share their electrons. This will result in the formation of a molecule of water (H2O). In this molecule, each hydrogen atom shares its electron with the oxygen atom, and the oxygen atom is sharing its electrons with the hydrogen atoms. Now, each atom has a stable valence shell. (Hydrogen follows the duet rule, meaning that it is stable with 2 valence electrons).

covalent bonding of water molecule

In covalent bonding, atoms can share one or more pairs of electrons to create double or even triple bonds, where multiple pairs of electrons are shared between atoms. This allows two of the same elements to covalently bond in different ratios to form different compounds. For example, Carbon, which has 4 valence electrons, and Oxygen which has 6 valence electrons can covalently bond to form either Carbon Monoxide (CO) or Carbon Dioxide (CO2).

In Carbon Monoxide, One Carbon and One Oxygen triple bond together and share 6 electrons between two atoms. (Each dash represents 1 pair of shared electrons).

carbon monoxide lewis structure

In Carbon Dioxide, One Carbon and Two Oxygens Bond together (Each dash represents 1 pair of shared electrons).

carbon dioxide lewis structure

QUICK TIP (Ionic vs. Covalent Bonding):

Ionic bonding occurs between a metal and a nonmetal.

Covalent bonding occurs only between nonmetals.

Determining Ionic or Covalent Bonding Quiz

That’ll be it for this article! As always, feel free to check us out on YouTube and Instagram at Biorithmetic! We hope you enjoyed reading and learning about ionic and covalent bonding. Make sure to try our available quiz below to check your understanding of this topic.

1. How many valence electrons does potassium (K) have?

a. 1

b. 2

c. 3

d. 4

2. How many valence electrons does nitrogen (N) have?

a. 1

b. 3

c. 5

d. 7

3. What type of bond will form between magnesium (Mg) and oxygen (O)?

a. Ionic

b. Covalent

4. What type of bond will form between carbon (C) and hydrogen (H)?

a. Ionic

b. Covalent

5. What type of bond will form between sodium (Na) and chlorine (Cl)?

a. Ionic

b. Covalent


1. a

2. c

3. a

4. b

5. a



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