Lewis Structures: Learn How to Draw Lewis Structures | Albert.io (2024)

Introduction to Lewis Structures

The only thing smaller than atoms is their subatomic particles; electrons, protons, and neutrons. Not even under a complex microscopic can we view the individual electrons that surround an atom’s nuclei. Lewis Dot Structures, or Lewis Structures for short, are visuals that represent the outermost shell of electrons, also known as valence electrons, and possible covalent bonds within an atom or molecule. These valence electrons are negatively charged and are attracted to the positively charged nucleus, made up of neutrons and protons. Keep in mind that in reality electrons are constantly moving around the nucleus and are not rooted in one place as portrayed in a 2D structure.

Lewis structures are drawn by a series of dots, lines, and atomic symbols and provide a structure for the way that the atom or molecule is arranged. A Lewis Dot Structure can be made for a single atom, a covalent compound, or a polyatomic ion.

Using the Periodic Table to Draw Lewis Dot Structures

The periodic table has all of the information needed to draw a Lewis dot structure. Each Group, or column, is indicated by a roman numeral which represents the number of valence electrons. For example, all elements which fall within the first column, or Group I, have one (1) valence electron. All elements in Group II have two (2) valence electrons, all the way up to VIII, eight (8) valence electrons. Properties are also consistent across the rows, or periods. Periods are indicated by a number, 1, 2, 3, etc. which represent the energy level, or shell of electrons. The first period, or row, has only one energy level that can hold a total of two electrons. Period 2, with a second shell, can hold a total of eight (8) electrons, also known as the octet rule. Period 3 and so forth can hold more than eight (8) electrons.

The periodic table also conveys electronegativity. The most electronegative elements are located in the uppermost right corner of the period table and decrease in electronegativity as you go down the Group or more left of a Period.

Throughout drawing Lewis dot structures, the periodic table will be a strong reference point when working with electrons, covalent bonding, and polyatomic ions.

Using Lewis Structures to Show Valence Electrons

Lewis dot structures can be drawn to show the valence electrons that surround an atom itself. This type of Lewis dot structure is represented by an atomic symbol and a series of dots. See the following examples for how to draw Lewis dot structures for common atoms involved in covalent bonding.

Example 1. Draw the Lewis Dot Structure for the Hydrogen atom.

Since Hydrogen is in Group I it has one (1) valence electron in its shell.

Example 2. Draw the Lewis Dot Structure for the Florine atom.

Since Fluorine is in Period 2, it can fit a maximum of eight (8) electrons second energy level. Fluorine Group VII, which means it has a total of seven (7) valence electrons around the atom.

Example 3. Draw the Lewis Dot Structure for Oxygen.

Since Oxygen is in Period 2, it can fit a maximum of eight (8) electrons second energy level. Oxygen Group VI, which means it has a total of six (6) valence electrons around the atom

Example A. Determine the total number of valence electrons for C

• Carbon is in Group IV, 4 valence electrons
• Total # of Valence Electrons in Carbon = 4

Example B. Determine the total number of valence electrons for H₂O

• Hydrogen, Group I, has 1 electron x 2 = 2
• Oxygen, Group VI, has 6 electrons x 1 = 6
• Total Valence Electrons in water = 8

Example C. Determine the total number of valence electrons for MgBr₂

• Magnesium, Group 2, has 2 electrons x 1 = 2
• Bromine, Group 7, has 7 electrons x 2 = 14
• Total # of Valence Electrons in MgBr₂ = 16

How to Draw a Lewis Structure

Step 1. Determine the total number of valence electrons to be depicted in the Lewis diagram.

Example: CO₂Total = 16

Step 2. Place least electronegative element in center and draw single bonds from the central atom to other atoms.

Step 3. Determine how many electrons must be added to the central element.

Assume that each outer element has a full valence (2 for H, 8 for everything else) from bonding and non-bonding electrons. Total all of these electrons, and subtract that from the total number of valence electrons in the molecule. CO₂ has 16 valence electrons. We assume each O has 8 valence electrons. 2×8=16; 16-16=0 therefore we don’t need to add any electrons to C

Step 4. Add double or triple bonds to the central atom until it has a full octet

Step 5. Add electrons to outer elements until they have full octets

Examples for Drawing Lewis Structures for Covalent Bonds

Here, we will be using the determined total number of valence electrons per atom and drawing them in the proper places. Reference the “How to Draw a Lewis Dot Structure” for a Step by Step guide. See the following Lewis dot structure diagrams for a few covalent compounds.

Example 1. Ammonia, NH₃

Nitrogen is in Group V which means it has a total of five (5) valence electrons. There are three (3) hydrogens present, each with their own sole electron giving the entire molecule a total of eight (8) to be accounted for. Since Nitrogen has 5 electrons and is looking for a total of 8 to fulfill its second energy shell, it is satisfied by the presence of 3 hydrogens which fulfills the octet rule. A nonbonding pair of Nitrogen is left and represented a pair of two dots.

Example 2. Methane, CH₄

In this structure, there are no nonbonding pairs of electrons present. All have been properly bonded in a series of lines representing two electrons each. In methane, each Hydrogen has the first energy shell filled with 2 electrons, its own valence electron along with a shared electron from Carbon, and Carbon’s second energy shell is filled with a total of eight electrons, 4 of its own and 4 shared (1 from each Hydrogen surrounding it).

Drawing Lewis Dot Structures for Polyatomic Ions

Lewis Dot Structures for drawing polyatomic ions are done very similarly to that of drawing individual atoms or covalent compounds. However, in this case, we will be dealing with ions rather than simple atoms. Ions are going to possess either a positive or negative charge which should be reflected by the number of electrons drawn as well as an indication of a “-“or “+”. This means that there are either additional electrons present to create a negative charge or less electrons present to create a more positive charge.

Example 1. PO₄^3-, Phosphate ion

With four oxygens present with six (6) electrons each and a phosphorus with five (5) there should be a total of 31 electrons. However, since there is a charge of -3.

There are a few things to keep in mind. Phosphorus is in Period 3, which means it can hold more than 8 electrons and creates a double bond to the oxygen which fulfills the octet rule for one oxygen, but not the others.

Example 2. NH₄⁺_, Ammonium ion

With ammonium, we are dealing with a positively charged polyatomic ion. The total valence electrons of the nitrogen and four hydrogens is 9 electrons. Since there is a positive charge of 1+ that means there is one less electron, so there will be a total of 8, which are represented by the four bonds as lines.

Example 3. OH^–, Hydroxide ion

The hydroxide ion has a total of how many electrons? Well, oxygen has 6 and hydrogen has 1, but since there is a negative charge on the ion, it will have an additional ion making a total 8 electrons, which are representing by the bonding pair between oxygen and hydrogen along with the 3 nonbonding (lone) pairs surrounding oxygen.

Key Concepts

Determine the total number of valence electrons of the element or compound. If a molecule has more than one element, add the valence electron of all elements present in the compound.

Determine which atom will be the central atom of the Lewis Dot Structure. The central atom is the least most electronegative atom in the compound. Remember the trend for electronegativity on the periodic table. Once determined, draw that element by atomic symbol in the center and draw single bonds to the other atoms.

Subtract full shell of valence electrons (2 for H, 8 for everything else) of each outer atom from the total number of valence electrons associated with the molecule. Distribute the remaining electrons to the central atom as non-bonding pairs

Form double and triple bonds until the central atom has a full octet.

Draw nonbonding pairs around the outer atoms until they have a full octet.

Check your work: Ensure that all of your valence electrons and bonds are accounted for.

Practice with Drawing Lewis Structures

1. Carbon
2. Sodium
3. Neon
4. HCl
5. H₂O
6. SO₂
7. NO₃^–
8. ClO₃^–
9. CN^–
10. SO₄^2-