Question

Use Lewis electron-dot symbols to represent the formation of (a) \(\mathrm{BrF}_{3}\) from bromine and fluorine atoms (b) AlF \(_{3}\) from aluminum and fluorine atoms

Short answer

BrF\textsubscript{3} has three single bonds and three lone pairs on Br. AlF\textsubscript{3} has three single bonds and no lone pairs on Al.

Step-by-step solution

Introduction to Lewis Structures

Lewis electron-dot symbols represent valence electrons around atoms. They help to visualize the formation of chemical bonds.

Determine the Total Valence Electrons for \textbf{BrF\textsubscript{3}}

Br has 7 valence electrons, and each F atom has 7 valence electrons. Total valence electrons = 7 + 3*7 = 28.

Draw the Br Atom as the Central Atom

Place Br in the center and surround it with 3 F atoms.

Distribute Electrons to Form Bonds and Fill Octets

Draw single bonds between Br and each F. Then place the remaining electrons to complete the octets.

Final Lewis Structure for \textbf{BrF\textsubscript{3}}

The structure has three single bonds with three lone pairs around Br and three lone pairs on each F atom.

Determine the Total Valence Electrons for \textbf{AlF\textsubscript{3}}

Al has 3 valence electrons, and each F atom has 7 valence electrons. Total = 3 + 3*7 = 24.

Draw the Al Atom as the Central Atom

Place Al in the center and surround it with 3 F atoms.

Distribute Electrons to Form Bonds and Fill Octets

Draw single bonds between Al and each F. Place the remaining electrons to complete the octets.

Final Lewis Structure for \textbf{AlF\textsubscript{3}}

The structure has three single bonds with no lone pairs around Al and three lone pairs on each F atom.

Key concepts

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. These electrons are crucial because they are primarily involved in chemical reactions and bond formations.
Understanding valence electrons helps us predict an element's reactivity and the types of bonds it can form.

To find the number of valence electrons in an atom, you can look at its group number on the periodic table. For example:

• Bromine (Br) is in group 17, so it has 7 valence electrons.
• Fluorine (F) is also in group 17, so it has 7 valence electrons.
• Aluminum (Al) is in group 13, so it has 3 valence electrons.

When forming compounds, atoms tend to share, lose, or gain valence electrons to achieve a stable electronic configuration.

In the case of \(\text{BrF}_3\):

• Br has 7 valence electrons, and each F has 7 valence electrons.
• Total valence electrons = 7 (Br) + 3*7 (F) = 28.

In \(\text{AlF}_3\):

• Al has 3 valence electrons, and each F has 7 valence electrons.
• Total valence electrons = 3 (Al) + 3*7 (F) = 24.

Chemical Bonds

Chemical bonds are the forces holding atoms together in a molecule. The two main types of chemical bonds are ionic and covalent.
In covalent bonding, atoms share pairs of electrons to attain a full outer shell.

For example, in \(\text{BrF}_3\) and \(\text{AlF}_3\):

• Each bromine (Br) atom shares one electron with each of the three fluorine (F) atoms to form three covalent bonds.
• Each aluminum (Al) atom shares one electron with each of the three fluorine (F) atoms, also forming three covalent bonds.

The shared electrons in these bonds allow the atoms involved to achieve a more stable configuration.

Visualizing these bonds with Lewis structures can clarify how electrons are shared or transferred around the atoms:

• For \(\text{BrF}_3\), draw single bonds between the Br atom and each of the three F atoms.
• For \(\text{AlF}_3\), draw single bonds between the Al atom and each of the three F atoms.

These structures help us see where the electrons are located and how they contribute to molecule's stability.

Octets

The octet rule states that atoms are most stable when they have eight electrons in their outermost shell. This rule guides how atoms bond with each other to form compounds.

Most elements seek to complete their outer shell with 8 electrons, resembling the electron configuration of noble gases, which are the most stable elements on the periodic table.
There are some exceptions to the octet rule, but many common elements follow this principle.

Let’s see this in action in \(\text{BrF}_3\) and \(\text{AlF}_3\):

• For \(\text{BrF}_3\), each Fluorine (F) atom achieves a full octet by sharing one electron with the central Bromine (Br) atom, which also ends up with more than 8 electrons, making it an exception.
• For \(\text{AlF}_3\), each Fluorine (F) atom achieves a full octet by sharing one electron with the central Aluminum (Al) atom. Aluminum has only 6 electrons in its outer shell—not a full octet, but it is still stable in this compound.

These examples show how atoms work towards achieving full outer shells through bonding, demonstrating the importance of the octet rule in chemical bonding.
Understanding this concept allows us to predict molecular structures and stability with greater accuracy.