Question

Challenge A nitrogen trifluoride molecule contains numerous lone pairs. Draw its Lewis structure.

Short answer

The Lewis structure for the NF3 molecule can be drawn by placing the Nitrogen (N) atom at the center and three Fluorine (F) atoms around it, forming single bonds between N and F atoms. Then, allocate the remaining 20 valence electrons as lone pairs: add three lone pairs to each Fluorine atom and one lone pair on Nitrogen. The final Lewis structure is: F | N - F | F with a lone pair on Nitrogen and three lone pairs on each of the Fluorine atoms.

Step-by-step solution

Determine valence electrons for each atom

First, we need to determine the number of valence electrons for each atom in the molecule. Nitrogen (N) is found in group 15/5A, which means it has 5 valence electrons. Fluorine (F) is found in group 17/7A, which means it has 7 valence electrons. Since there are three fluorine atoms, we will have a total of 21 valence electrons from all three fluorine atoms.

Calculate the total number of valence electrons in the molecule

Now, we need to calculate the total number of valence electrons for the whole NF3 molecule. To do this, simply add the number of valence electrons from nitrogen and the total valence electrons from all fluorine atoms together: Total valence electrons = Valence electrons of Nitrogen + Valence electrons of 3 Fluorine atoms Total valence electrons = 5 (from N) + 21 (from 3F) Total valence electrons = 26

Draw the Lewis structure

To form the Lewis structure of the NF3 molecule, we will start by placing the Nitrogen atom at the center since it has a lower electronegativity than the Fluorine atoms. This means that Nitrogen tends to form bonds with Fluorine atoms more easily. Then, place the three Fluorine atoms around the Nitrogen atom. 1. Draw single bonds between the central Nitrogen atom and each of the Fluorine atoms, using up 2 electrons for each bond. This will require 6 electrons in total (3 bonds × 2 electrons per bond). 2. Subtract the electrons used in bonding (6) from the total valence electrons (26). This gives us 20 electrons remaining to be distributed as lone pairs. 3. Now, place the remaining electrons as lone pairs around each fluorine atom, with three lone pairs per Fluorine atom, giving each atom an octet structure. This will require 18 electrons (3 lone pairs × 3 Fluorine atoms × 2 electrons per lone pair). 4. There are two electrons left which we are going to place as a lone pair on the central Nitrogen atom. So, the final Lewis structure for the NF3 molecule will look like this: F | N - F | F with a lone pair on Nitrogen and three lone pairs on each of the fluorine atoms.

Key concepts

Valence Electrons

Understanding valence electrons is key to mastering Lewis structures. Valence electrons are the outermost electrons of an atom. They play a crucial role in chemical bonding because they are the ones involved in forming bonds.
Each element in the periodic table has a specific number of valence electrons. For instance:

• Nitrogen (N), found in group 15 (or 5A) of the periodic table, has 5 valence electrons.
• Fluorine (F), in group 17 (or 7A), possesses 7 valence electrons.

Knowing the number of valence electrons allows you to predict how atoms will bond in a molecule. It will help you accurately draw Lewis structures, illustrating how atoms share or transfer electrons.

Nitrogen Trifluoride

Nitrogen trifluoride (NF extsubscript{3}) is a chemical compound composed of one nitrogen atom and three fluorine atoms. It's a fascinating molecule to study because of its subtle complexity in terms of chemical bonding and lone pairs.
In the NF extsubscript{3} molecule, nitrogen acts as the central atom. It connects to the three surrounding fluorine atoms. This arrangement is typical because nitrogen can effectively share its electrons with highly electronegative atoms like fluorine.
As a result, this molecule displays a trigonal pyramidal shape due to the distribution of electrons and lone pairs. This geometric detail is crucial for understanding many of its chemical properties.

Chemical Bonding

At the heart of understanding Lewis structures is the concept of chemical bonding. Chemical bonds form when atoms share or transfer valence electrons to achieve a full outer electron shell, usually with eight electrons. This is often referred to as following the "octet rule."
In nitrogen trifluoride:

• Each nitrogen-fluorine bond is a single bond, comprised of two shared electrons.
• Nitrogen forms three single bonds with three fluorine atoms, utilizing 6 of the available valence electrons, creating a stable configuration for the molecule.

The stability of NF extsubscript{3} is due to the sharing of electrons between nitrogen and fluorine atoms, fulfilling their valence shell requirements.

Lone Pairs

Lone pairs are pairs of valence electrons that are not shared with another atom in a bond. In a Lewis structure, they are specifically highlighted because they affect the shape and properties of a molecule.
For nitrogen trifluoride (NF extsubscript{3}):

• Each fluorine atom holds three lone pairs, which account for 18 electrons and complete the valence shell, giving each fluorine an octet.
• Nitrogen also has one lone pair contributing to its structural shaping and affects the molecule's geometry.

Lone pairs are significant because they don't just fill spaces; they influence molecular shapes and reactivities by repelling bonding pairs, altering angles and creating distinct molecular geometries.