Animated Solution for Chemistry - Chemical Bonding and Molecular Structure: The compound(s) with TWO lone pairs of electrons on the central atom is(are)

Have you ever looked at a chemical formula and wondered what invisible forces are shaping its structure? Molecules aren't just random clusters of atoms; they are highly organized, three-dimensional architectures dictated by the delicate dance of electrons. Today, we are embarking on a thrilling quest to uncover the hidden lone pairs in four fascinating molecules: ${\text{BrF}}_5$, ${\text{ClF}}_3$, ${\text{XeF}}_4$, and ${\text{SF}}_4$. Our mission? To find out which of these central atoms are secretly hoarding exactly two lone pairs of electrons.

Steric Number

To crack this puzzle, we don't need to guess. We have a powerful mathematical tool at our disposal: the Steric Number ($H$). This number tells us the total number of electron domains (both bonds and lone pairs) around a central atom.

The formula is beautifully simple:

Here, $V$ represents the valence electrons of the central atom, and $M$ is the number of monovalent surrounding atoms (like halogens or hydrogen). If we had a charge, we would subtract the cation charge ($C$) or add the anion charge ($A$), but our molecules today are all neutral. Once we have the steric number, finding the lone pairs ($LP$) is just a matter of subtraction:

Let's unleash this formula on our contenders!

Option A: Bromine Pentafluoride (${\text{BrF}}_5$) Imagine Bromine sitting at the center. As a proud member of the halogen family (Group 17), it brings 7 valence electrons to the table. It is surrounded by 5 monovalent fluorine atoms.

Plugging this into our master equation:

A steric number of 6 corresponds to an $s{p}^3{d}^2$ hybridization. Since there are 5 fluorine atoms, we have 5 bond pairs.

Bromine has just one lone pair, giving the molecule a square pyramidal shape. Close, but not what we are looking for!

Option B: Chlorine Trifluoride (${\text{ClF}}_3$) Next up is Chlorine, another halogen with 7 valence electrons. This time, it's bonded to only 3 fluorine atoms.

A steric number of 5 means $s{p}^{3}d$ hybridization. Subtracting the 3 bond pairs:

Bingo! We found our first winner. Those two lone pairs occupy the spacious equatorial positions to minimize repulsion, forcing the molecule into a distinctive T-shape.

Option C: Xenon Tetrafluoride (${\text{XeF}}_4$) Now, let's look at Xenon. Being a noble gas, it usually likes to keep to itself with a full octet of 8 valence electrons. But fluorine is highly electronegative and forces Xenon to bond.

With a steric number of 6 and 4 bond pairs:

Another perfect match! Xenon holds onto two lone pairs. To stay as far away from each other as possible, these lone pairs sit opposite each other (axially), leaving the four fluorine atoms in a beautiful square planar arrangement.

Option D: Sulfur Tetrafluoride (${\text{SF}}_4$) Finally, we examine Sulfur from Group 16, bringing 6 valence electrons. It bonds with 4 fluorine atoms.

Subtracting the 4 bond pairs:

Sulfur has only one lone pair, which takes an equatorial spot, creating a playful see-saw geometry. Not our target today.

Through the elegant logic of the steric number, we've unmasked the hidden electron pairs. Both ${\text{ClF}}_3$ and ${\text{XeF}}_4$ possess exactly two lone pairs on their central atoms. The invisible forces of electron repulsion dictate their unique shapes, proving once again that in chemistry, what you can't see is just as important as what you can. Keep exploring, keep questioning, and never underestimate the power of a lone pair!