Question

Write Lewis structures for \(\mathrm{BrF}_{3}, \mathrm{ClF}_{5},\) and \(\mathrm{IF}_{7}\). Identify those in which the octet rule is not obeyed.

Short answer

None obey the octet rule; all exceed it.

Step-by-step solution

Understanding Lewis Structures

A Lewis structure is a diagram representing the arrangement of valence electrons around atoms in a molecule. Lewis structures help us understand the bonding and the distribution of electrons around atoms within a molecule.

Draw Lewis Structure for \(\mathrm{BrF}_3\)

1. Count the valence electrons: Bromine (Br) has 7 valence electrons, and each Fluorine (F) has 7 valence electrons. Thus, BrF₃ has \(7 + 3(7) = 28\) electrons. 2. Place Bromine in the center and draw single bonds to three Fluorine atoms. Each bond uses 2 electrons, totaling 6 electrons.3. Subtracting these 6 electrons from 28 leaves 22 electrons.4. Place 6 electrons on each Fluorine to complete their octet, using up 18 electrons.5. There are now 4 electrons remaining, which become two lone pairs on the Bromine.6. The structure of BrF₃ is: three bonded pairs and two lone pairs, making it a T-shaped geometry. This does not obey the octet rule on the central atom since Br has 10 electrons.

Draw Lewis Structure for \(\mathrm{ClF}_5\)

1. Count the valence electrons: Chlorine (Cl) has 7 valence electrons, and each Fluorine (F) has 7 valence electrons. Thus, ClF₅ has \(7 + 5(7) = 42\) electrons.2. Place Chlorine at the center and form single bonds with five Fluorine atoms, using 10 electrons.3. After subtracting these, 32 electrons remain.4. Distribute the remaining electrons as lone pairs to satisfy the octet of all five Fluorine atoms, using 30 electrons.5. Two electrons remain, forming one lone pair on the Chlorine.6. ClF₅ forms a square pyramidal geometry. Here, chlorine exceeds the octet rule with 12 electrons.

Draw Lewis Structure for \(\mathrm{IF}_7\)

1. Count the valence electrons: Iodine (I) has 7 valence electrons, and each Fluorine (F) has 7 valence electrons. Thus, IF₇ has \(7 + 7(7) = 56\) electrons.2. Place Iodine at the center and link it to seven Fluorine atoms using single bonds, which requires 14 electrons.3. After this, 42 electrons remain.4. Distribute the rest as lone pairs around the seven Fluorine atoms, using up all 42 electrons.5. The resulting structure is that of a pentagonal bipyramidal arrangement.6. Here, iodine does not conform to the octet rule, possessing 14 electrons.

Identify Octet Rule Violations

In these molecules, \(\mathrm{BrF}_3\), \(\mathrm{ClF}_5\), and \(\mathrm{IF}_7\), the central halogen atoms do not obey the octet rule. Specifically, Br, Cl, and I exceed eight valence electrons with 10, 12, and 14 electrons respectively.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a significant role in chemical bonding. They are crucial because they participate in forming bonds with other atoms. Knowing the number of valence electrons helps predict how atoms will interact and bond with each other.
For atoms in Group 17 of the periodic table, such as bromine, chlorine, and iodine, there are 7 valence electrons. When atoms form molecules, these electrons are shared or exchanged to help achieve greater stability.
Understanding valence electrons is essential when drawing Lewis structures, as it determines how many electrons are available for bonding and how they can be distributed around the atoms in the structure.

Octet Rule

The octet rule is a chemical guideline stating that atoms are most stable when they have eight electrons in their valence shell. This rule is inspired by the electronic configuration of noble gases, which have full outer electron shells.
However, the octet rule does have its exceptions, especially in molecules where the central atom can have more than eight electrons. This is often observed in elements from Period 3 and beyond, such as sulfur, phosphorus, chlorine, and iodine.
In some compounds like BrF₃, ClF₅, and IF₇, the central halogen atoms possess more than eight electrons, showing a common exception to the octet rule in larger molecules.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. This arrangement can affect a molecule's physical and chemical properties, including reactivity, polarity, and biological activity.
Understanding molecular geometry helps predict how molecules will interact with each other and their environment. Factors like the number of bonding pairs and lone pairs around the central atom influence the final shape of the molecule.
Analyzing molecular geometry often involves considering repulsions between electron pairs, as described by the VSEPR (Valence Shell Electron Pair Repulsion) theory, which helps determine the most favorable structure.

Bromine Trifluoride (BrF3)

Bromine trifluoride (BrF₃) is an interhalogen compound where bromine is bonded to three fluorine atoms. It has a total of 28 valence electrons. Of these, 6 are used to form three Br-F bonds, leaving bromine with two lone pairs.
The molecule takes on a T-shaped geometry according to VSEPR theory because of the repulsion between lone pairs and bond pairs. In this structure, bromine is surrounded by a total of 10 electrons, exceeding the octet rule. This arrangement arises due to bromine's ability to expand its valence shell and accommodate more electrons.

Chlorine Pentafluoride (ClF5)

Chlorine pentafluoride (ClF₅) is a compound where chlorine is centrally bonded to five fluorine atoms, involving 42 valence electrons in total. Ten of these electrons create Cl-F bonds, while the remaining electrons form lone pairs.
The geometry of ClF₅ is square pyramidal, where the five fluorine atoms are positioned around chlorine in a square planar shape, with an additional lone pair on chlorine, influencing the molecule's shape.
Chlorine in ClF₅ has 12 electrons in its valence shell, surpassing the limits of the octet rule, illustrating how elements in and beyond the third period can hold more than eight electrons due to available d orbitals.

Iodine Heptafluoride (IF7)

Iodine heptafluoride (IF₇) is a compound featuring iodine bonded to seven fluorine atoms. The molecule has a total of 56 valence electrons. Fourteen electrons are used to form seven I-F bonds, with the remaining distributed as lone pairs on the fluorine atoms.
The molecular geometry of IF₇ is described as pentagonal bipyramidal. This structure places five fluorine atoms in a plane with two axial fluorines positioned above and below this plane.
In IF₇, iodine accommodates 14 electrons, well beyond the traditional octet, once again highlighting the ability of larger atoms like iodine to expand their valence shells by utilizing available d orbitals.