VSEPR for 5 electron clouds (part 1) | AP Chemistry | Khan Academy

Summary

This video explains how to use VSEPR theory to predict the molecular shapes of phosphorus pentachloride (PCl₅) and sulfur tetrafluoride (SF₄). It begins by drawing Lewis structures to represent valence electrons, followed by identifying electron clouds around the central atoms. Using VSEPR theory, it demonstrates that PCl₅ adopts a trigonal bipyramidal shape, while SF₄ forms a seesaw shape due to the lone pair on sulfur. Bond angles and electron pair repulsion are discussed to explain the stability of these molecular geometries.

Takeaways

• 🧪 **VSEPR Theory Application**: The script uses VSEPR theory to predict molecular geometry, focusing on electron pair repulsion.
• 🌐 **Dot Structure Creation**: It explains how to draw a dot structure for molecules, showing valence electrons and their distribution.
• 📊 **Electron Counting**: The process of counting valence electrons for phosphorus pentachloride and sulfur tetrafluoride is detailed.
• 🔬 **Central Atom Placement**: Phosphorus is placed at the center of its molecule due to chlorine's higher electronegativity, similarly sulfur for its compound.
• 📐 **Electron Clouds Identification**: The script identifies electron clouds around the central atoms and their influence on molecular shape.
• 🔄 **Repulsion and Geometry**: It discusses how electron pairs repel each other to achieve the most stable molecular geometry.
• 📈 **Trigonal Bipyramidal Shape**: For molecules with five electron clouds, a trigonal bipyramidal shape is predicted by VSEPR theory.
• 📏 **Bond Angles**: The script explains the ideal bond angles in a trigonal bipyramidal geometry: 120° for equatorial, 90° for axial-equatorial, and 180° for axial-axial.
• 🧩 **Lone Pairs Impact**: The significance of lone pairs in determining molecular shape and their repulsion effects are highlighted.
• 🔍 **Formal Charge Consideration**: Assigning formal charges helps in understanding the stability and structure of molecules.
• 🛠️ **Molecular Shape Prediction**: The final shape of a molecule is predicted by ignoring lone pairs and considering the positions of bonding atoms.

Q & A

What is the first step in predicting the structure of phosphorus pentachloride (PCl5) using VSEPR theory?

The first step is to draw the Lewis dot structure to show the valence electrons. Phosphorus has 5 valence electrons, and each of the five chlorine atoms has 7 valence electrons, giving a total of 40 valence electrons.

Why does phosphorus exceed the octet rule in phosphorus pentachloride (PCl5)?

Phosphorus exceeds the octet rule because it is in Period 3 of the periodic table. Atoms in Period 3 and beyond can expand their valence shells to accommodate more than 8 electrons.

How many electron clouds surround the central phosphorus atom in phosphorus pentachloride (PCl5), and what is their arrangement?

There are five electron clouds around the central phosphorus atom, corresponding to the five P-Cl bonds. These electron clouds are arranged in a trigonal bipyramidal shape.

What are the bond angles in the trigonal bipyramidal structure of phosphorus pentachloride (PCl5)?

In the trigonal bipyramidal structure, the bond angles between the three equatorial chlorines are 120°, the bond angle between the two axial chlorines is 180°, and the bond angles between the axial and equatorial chlorines are 90°.

What determines the geometry of the electron clouds in VSEPR theory?

The geometry of the electron clouds is determined by the repulsion between negatively charged valence electrons. These electron clouds arrange themselves to be as far apart as possible, minimizing repulsion.

Why is sulfur tetrafluoride (SF4) considered to have a lone pair of electrons?

Sulfur in SF4 has 6 valence electrons, and after forming bonds with four fluorine atoms, two electrons remain as a lone pair on the sulfur atom, exceeding the octet rule due to sulfur being in Period 3.

How many electron clouds are around the central sulfur atom in sulfur tetrafluoride (SF4), and how do they arrange themselves?

There are five electron clouds around the central sulfur atom in SF4 (four bonding pairs and one lone pair). These electron clouds arrange themselves in a trigonal bipyramidal geometry.

Why is the lone pair of electrons placed in the equatorial position in sulfur tetrafluoride (SF4)?

The lone pair is placed in the equatorial position to minimize repulsion. Lone pairs take up more space than bonding pairs, and placing it in the equatorial position results in fewer 90° interactions, reducing electron pair repulsion.

What is the molecular geometry of sulfur tetrafluoride (SF4) after considering VSEPR theory?

The molecular geometry of sulfur tetrafluoride (SF4) is a seesaw shape because the lone pair of electrons distorts the trigonal bipyramidal arrangement.

How does the seesaw shape of sulfur tetrafluoride (SF4) relate to the VSEPR model?

The seesaw shape arises because the lone pair of electrons in SF4 causes greater repulsion and distorts the ideal trigonal bipyramidal geometry. By ignoring the lone pair, the remaining bonds form a seesaw-like structure.

Outlines

00:00

🧪 VSEPR Theory for Phosphorus Pentachloride

This paragraph explains the use of VSEPR theory to determine the structure of phosphorus pentachloride (PCl₅). It begins by constructing a Lewis dot structure for the molecule, where phosphorus is the central atom, surrounded by five chlorine atoms. The total valence electron count is calculated as 40, which is distributed according to the octet rule for chlorine and an expanded octet for phosphorus. The shape predicted by VSEPR theory is a trigonal bipyramidal configuration due to five electron clouds surrounding the central phosphorus atom. The description covers bond angles of 120° for equatorial chlorines, 180° between axial chlorines, and 90° between axial and equatorial chlorines.

05:02

🧬 Trigonal Bipyramidal Geometry Explained

This paragraph emphasizes the importance of understanding trigonal bipyramidal geometry, which is commonly adopted by molecules with five electron clouds. The paragraph highlights that molecules with this configuration tend to have electron clouds arranged in equatorial and axial positions, with three electron clouds equatorially and two axially. In the absence of lone pairs, the molecule takes on this trigonal bipyramidal shape. The example of sulfur tetrafluoride (SF₄) is introduced, and a step-by-step process is outlined for drawing its Lewis dot structure, taking into account sulfur's lone pair and fluorine's octet. The electron clouds are arranged to minimize repulsion, and different arrangements of the lone pair in equatorial and axial positions are compared.

10:02

🔎 Minimizing Lone Pair Repulsion in Trigonal Bipyramids

The third paragraph delves into the analysis of lone pair repulsion in trigonal bipyramidal structures. Using SF₄ as an example, it compares two configurations: lone pairs in the equatorial versus axial positions. The paragraph explains that lone pairs in equatorial positions cause less repulsion than in axial positions because they are only 90° to two bonding pairs rather than 90° to three, as seen in the axial arrangement. Therefore, VSEPR theory predicts the equatorial placement as more stable. Finally, the structure is confirmed to be a seesaw shape due to the lone pair's influence, and the paragraph uses the seesaw analogy to explain the resulting geometry.

Keywords

💡VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) Theory is used to predict the 3D structure of molecules based on the repulsion between electron pairs in the valence shell. In the video, VSEPR is applied to predict the shape of phosphorus pentachloride (PCl5) and sulfur tetrafluoride (SF4). The theory helps explain why electron clouds arrange themselves to minimize repulsion, leading to specific molecular geometries.

💡Trigonal Bipyramidal

A trigonal bipyramidal shape is a molecular geometry where five electron clouds arrange themselves with three atoms in a planar triangular arrangement (equatorial) and two atoms in axial positions. This structure is used in the video to describe the shape of phosphorus pentachloride (PCl5), where three chlorine atoms are in a plane, and two are above and below that plane.

💡Electron Clouds

Electron clouds refer to regions of electron density around an atom, which include bonding and non-bonding (lone) pairs of electrons. In the video, electron clouds are central to determining molecular shapes via VSEPR theory, as they repel each other and spread out to minimize repulsion, affecting the overall geometry of the molecule.

💡Phosphorus Pentachloride (PCl5)

Phosphorus pentachloride (PCl5) is a chemical compound discussed in the video to demonstrate VSEPR theory. It has five chlorine atoms bonded to a central phosphorus atom. The structure adopts a trigonal bipyramidal shape, with bond angles of 120° and 90° between chlorine atoms due to electron cloud repulsion.

💡Formal Charge

Formal charge is a concept used to calculate the charge on individual atoms in a molecule. It helps in determining the most stable Lewis structure. In the video, formal charge is mentioned to explain why phosphorus in PCl5 has a charge of zero, which helps validate the drawn dot structure.

💡Octet Rule

The octet rule states that atoms tend to bond in such a way that they have eight electrons in their valence shell, achieving stability similar to noble gases. In the video, the rule is applied to chlorine atoms in PCl5, ensuring each chlorine atom is surrounded by eight electrons. However, phosphorus exceeds the octet rule due to its position in Period 3.

💡Expanded Valence Shell

An expanded valence shell refers to elements in Period 3 or beyond that can accommodate more than eight electrons in their valence shell. In the video, phosphorus in PCl5 and sulfur in SF4 are examples of atoms that exceed the octet rule, as they can hold more than eight electrons due to available d-orbitals.

💡Equatorial Positions

Equatorial positions refer to the three atoms that lie in the same plane in a trigonal bipyramidal geometry. In the video, the three chlorine atoms in PCl5 and the fluorines in SF4's trigonal bipyramidal structure occupy these positions to minimize repulsion, creating 120° bond angles.

💡Axial Positions

Axial positions are the two atoms located above and below the plane in a trigonal bipyramidal shape. In the video, the axial chlorine atoms in PCl5 and SF4 are discussed in relation to the bond angles (180° between them) and how these positions interact with equatorial atoms, influencing molecular shape.

💡Seesaw Shape

A seesaw shape occurs when there are five regions of electron density but one is a lone pair, leading to a distorted trigonal bipyramidal shape. In the video, sulfur tetrafluoride (SF4) is shown to adopt this shape, where the lone pair takes one of the equatorial positions, causing a 'seesaw' geometry with fluorine atoms.

Highlights

Phosphorus pentachloride (PCl5) has 40 valence electrons: 5 from phosphorus and 35 from chlorine.

Phosphorus is placed in the center of the dot structure as it's less electronegative than chlorine.

Chlorine atoms each need 6 more electrons to complete their octet, leading to a total of 30 valence electrons distributed.

Phosphorus can exceed the octet rule because it is in Period 3, allowing for expanded valence shells.

VSEPR theory is used to predict the molecular shape by counting electron clouds, including both bonding and non-bonding electron regions.

Phosphorus pentachloride has five electron clouds around the central atom, leading to a trigonal bipyramidal shape.

Three chlorine atoms are in equatorial positions, with bond angles of 120°, while two are in axial positions with a 180° bond angle between them.

The bond angle between axial and equatorial positions is 90°, resulting in a mix of 120°, 180°, and 90° angles.

Sulfur tetrafluoride (SF4) has a total of 34 valence electrons: 6 from sulfur and 28 from fluorine.

Sulfur also exceeds the octet rule, and the structure includes a lone pair of electrons on sulfur.

Sulfur tetrafluoride adopts a trigonal bipyramidal shape due to the five electron clouds around sulfur.

In SF4, lone pairs take up more space and are placed in equatorial positions to minimize repulsion in a trigonal bipyramidal structure.

VSEPR predicts that the structure with the lone pair in the equatorial position is more stable due to less electron repulsion.

Sulfur tetrafluoride has a seesaw molecular shape, with two axial and two equatorial fluorine atoms, after ignoring the lone pair.

The seesaw shape is due to the asymmetric distribution of electrons and the lone pair repelling bonding pairs more effectively in the equatorial plane.

Transcripts

Let's use VSEPR theory to predict

the structure of this molecule-- so phosphorus pentachloride.

So the first thing we need to do is draw a dot structure

to show our valence electrons.

We find phosphorus in Group 5.

So 5 valence electrons.

Chlorine in Group 7.

So 7 valence electrons, and I have 5 of them.

So 7 times 5 is 35.

Plus 5 gives us a total of 40 valence electrons

that we need to show in our dot structure.

So phosphorus goes in the center because it is not

as electronegative as chlorine.

And we have five chlorines.

So we go ahead and put our five chlorines

around our central phosphorus atoms like that.

If we see how many valence electrons we've drawn so far,

this would to be 2, 4, 6, 8, and 10.

So 40 minus 10 gives us 30 valence electrons left over.

And remember, you start putting those leftover electrons

on your terminal atom.

So we're going to put those on the chlorines.

Each chlorine's going to follow the octet rule.

So that means each chlorine needs 6 more electrons.

Now, each chlorine is surrounded by 8 valence electrons

like that.

So if I'm adding 6 more electrons to 5 atoms, 6 times 5

is 30.

So I have now represented all of my valence electrons

on my dot structure.

Notice that phosphorus is exceeding the octet rule here.

There are 10 valence electrons around phosphorus.

And it's OK for phosphorus to do that because it's in Period 3

on the periodic table.

I like to think about formal charge.

And so if you assign a formal charge to phosphorus,

you'll see it has a formal charge of 0.

And that helps to explain-- for me, anyway-- the resulting dot

structure.

Now, step two.

We're going to count the number of electron clouds

that surround our central atom.

Remember, an electron cloud is just

a region of electron density.

So I could think about these bonding electrons in here

as a region of electron density around my central atom.

I could think about these bonding electrons, too.

So here's another electron cloud.

And you can see we have a total of five electron

clouds around our central atom.

The next step is to predict the geometry of the electron

clouds.

Those valence shell electrons are going to repel each other.

All right.

So that's a VSEPR theory-- Valence Shell Electron Pair

Repulsion.

Since they're all negatively charged,

they're going to repel and try to get as far

away from each other as they possibly can in space.

When you have five electron pairs,

it turns out the furthest they can get away from each other

in space is a shape called a trigonal bipyramidal shape.

So let me see if I can draw our molecule in that shape.

We're going to have our phosphorus in the center,

and we're going to have three chlorines on the same plane.

So let me attempt to show three chlorines on the same plane

here.

These are called the equatorial positions

because they're kind of along the equator, if you will.

So three chlorines in the same plane, one chlorine

above the plane, and one chlorine below the plane.

Those are called axial positions.

All right.

So there's a quick sketch.

Let me see if I can draw a slightly better shape

of a trigonal bipyramidal shape here.

So let me see if I can draw one over here so you

can see what it looks like a little bit better.

So we could have one pyramid looking something like that.

And then, down here, let's see if we

can draw another pyramid in here like that.

So that's a rough drawing, but we're

trying to go for a trigonal bipyramidal shape here.

So let's focus in on those chlorines that

are on the same plane first.

If I'm looking at these three chlorines

and I go over here to my trigonal bipyramidal shape,

you could think about those three chlorines

as being at these corners here.

So it's a little bit easier to see.

They're in the same plane.

So those are the equatorial chlorines.

When I think about the bond angle for those--

so those chlorines being in the same plane,

you have these three bond angles here.

And so when we did trigonal planar,

we talked about 360 degrees divided by 3--

giving us a bond angle of 120 degrees.

So you could think about that as being a bond angle of 120.

All right.

So same idea.

Those bonding electrons are going to repel each other.

When we focus in on our axial chlorines-- so this one up here

and this one down here.

You could think about those as being here and here

on your trigonal bipyramidal shape like that.

And if you draw the axis, if you draw a line down this way

connecting those, it's easy to see

those are 180 degrees from each other.

So you could think about a bond angle

of 180 degrees between your chlorines like that.

And then, finally, if we think about the bond angle between,

let's say, this axial chlorine up here at the top

and then one of these green chlorines right here,

I think it's a little bit easier to see that's 90 degrees here.

So this bond angle right here would be 90 degrees.

And so those are your three ideal bond angles

for a trigonal bipyramidal situation here.

It's important to understand this trigonal bipyramidal shape

because all of the five electron cloud drawings that we're

going to do are going to have the electron clouds

want to take this shape.

So it's important to understand those positions.

For step four, ignore any lone pairs

and predict the geometry of the molecule.

Well, there are no lone pairs on our central phosphorus.

So the electron clouds take a trigonal bipyramidal shape

and so does the molecule.

Let's go ahead and do another example.

Sulfur tetrafluoride here.

So we're going to start by drawing the dot structure,

and we need to count our valence electrons, of course.

So sulfur's in Group 6, so 6 valence electrons.

Fluorine is in Group 7.

So 7 valence electrons.

I have 4 of them.

7 times 4 is 28.

28 plus 6 is 34 valence electrons.

We know sulfur is going to go in the center

because fluorine is much more electronegative.

We put sulfur in the center here.

We know sulfur is bonded to 4 fluorines.

So we put our fluorines around like that.

And let's see how many valence electrons we've shown so far--

2, 4, 6, and 8.

So 34 minus 8 gives us 26 valence electrons

we still need to account for on our dot structure.

We're going to start by putting those leftover

electrons on our terminal atoms, which are our fluorines.

Fluorine's going to have an octet of electrons around it.

Therefore, each fluorine needs six,

since each fluorine already has two around it.

So we go ahead and put 6 valence electrons

around each one of our fluorine atoms.

All right.

So we are showing 6 more valence electrons on 4 atoms.

6 times 4 is 24.

So 26 minus 24 gives us 2 leftover valence electrons.

And remember your rules for drawing dot structures.

When you get some leftover electrons,

you're going to go ahead and put them on your central atom now.

So we have a lone pair of electrons on our sulfur.

And by adding that lone pair of electrons to our sulfur,

the sulfur now exceeds the octet rule.

But once again, it's OK for sulfur

to have an expanded valence shell.

It's in Period 3 on our periodic table.

And once again, I like to think about formal charge.

And if you assign a formal charge to that sulfur,

it has a formal charge of 0.

So that just helps me understand these dot structures

a little bit better.

So we've drawn our dot structure.

Let's go back up and remind ourselves of the next step

here.

So once you complete step one, next

is the electron cloud step.

So how many electron clouds do you

have surrounding your central atom?

So we go back down, and we look at our electron clouds

that surround our central atom.

Here's the regions of electron density.

So we know that these bombing electrons here,

that would be one electron cloud.

Same with these bonding electrons.

Same with these bonding electrons.

And same with these bonding electrons.

And then we have a lone pair of electrons on our sulfur.

Well, that's also a region of electron density surrounding

our central atom.

So that lone pair you could think of as being an electron

cloud as well.

And so we have five electron clouds.

So just like in the previous example.

And when you have five electron clouds,

those electron clouds are going to try

to adopt a trigonal bipyramidal shape-- just like we saw,

again, in the previous example.

So let's go ahead and draw two possible versions of the dot

structure for this molecule.

All right.

So I'm going to draw one right here.

And for this first version, I'm going

to show the lone pair of electrons

on the sulfur in the equatorial position here.

So I'm going to put the lone pair of electrons right here,

equatorial here.

And so that means there are two fluorines also equatorial.

And then that means that there's one fluorine here, axial.

Another fluorine here, axial.

So that is one possible dot structure.

The other possibility would be, of course,

to put to the lone pair of electrons

in the axial position.

So if we do that, we would have a sulfur bonded

to three fluorines.

Those would be the equatorial fluorines.

And then, we would have a lone pair of electrons.

Let's just put it right here in the axial position.

And then, another fluorine in the axial position.

So here are our two possibilities.

So let's see if we can analyze this structure.

Now, when you have lone pairs of electrons in your dot

structure, lone pairs take up more space.

Or non-bonding electrons-- I should say-- take up

more space than bonding electrons.

And so since they take up more space,

they're going to repel a little bit more.

And so that means that when you're

trying to figure out valence shell electron pair repulsion,

the lone pairs of electrons are more important to focus on

in terms of where you're going to put them.

Let's focus in on those lone pairs of electrons,

and let's think about how they're

going to repel the other electrons in these two dot

structures.

Let's look at the left here where

we had the lone pair of electrons

in the equatorial position here.

And if you're thinking about how they're interacting

with, let's say, these bonding electrons in the same plane

here, this is about 120 degrees between the bonding electrons

and the non-bonding electrons.

And it turns out that 120 degrees is not

as important in terms of repelling as, say, something

like 9 degrees.

You tend to ignore the 120 degree interactions when

you're analyzing these structures.

However, a 90-degree angle between a bonding pair

and a non-bonding pair-- and we had that example

for-- let me go ahead and show you right here.

So let's think about this lone pair of electrons repelling

these bonding electrons.

So in the axial position.

Well, these two are only 90 degrees away.

So remember, 9 degrees, of course, being closer,

you're going to get more repulsion from this interaction

than in the previous interactions.

So we're going to focus in on the 90-degree interactions

here.

Those bonding electrons and non-bonding electrons

repel each other.

And you have one possibility with the axial fluorine.

You also have another possibility

with this axial fluorine.

Essentially, you have a lone pair of electrons, 90 degrees,

from two pairs of bonding electrons

from the example on the left.

And of course, that's going to destabilize it somewhat.

But let's compare this dot structure

with the one on the right now.

So we have our lone pair of electrons in the axial position

this time.

And you can see that we have three fluorines

in the equatorial positions.

So you have these bonding electrons

in the equatorial position, which

means that that lone pair of electrons

is 90 degrees to all three of those.

And so that, of course, is going to cause

some serious repulsion, so 90 degrees to 3.

In the example on the right, you have these three interactions--

90 degrees.

An example on the left, you have only two of these.

The goal, of course, is to minimize electron pair

repulsion.

So VSEPR theory actually predicts

that this dot structure on the left is the correct one.

You're going to see-- in the next video--

that non-bonding electrons are placed

in equatorial positions in trigonal bipyramids

to minimize electron pair repulsion.

So just think about putting your lone pairs of electrons

in the equatorial position.

So the structure on the left wins.

Let's go ahead and redraw that so we

can analyze it a little bit better.

All right.

So I have my sulfur in the center here.

Let's go ahead and change colors.

So I'm going to put the sulfur in the center.

I have my fluorine in a plane.

Another fluorine in the plane.

My lone pair of electrons in a plane.

All right.

Those are my equatorial ones.

I have a fluorine this way.

And I have a fluorine that way.

So when you're looking at bond angles, of course,

between this sulfur fluorine bond angle--

the ideal bond angle anyway-- would be 120 degrees.

So we can say and we would expect it to be 120 degrees.

If you're talking about this axial fluorine

and this equatorial one, we would

expect that to be 90 degrees.

And then, finally, between the two axial

fluorines-- so this bond angle back here

would, of course, be 180 degrees.

OK.

So we've done a lot of talking, and we still haven't even

talked about the final name for the shape of this molecule.

So let's go back up here and look at our rules really fast.

So we've done a lot of work to predict

the geometry of the electron clouds around the central atom

and draw it.

And finally, we get to predict the shape of the molecule.

And we do that by ignoring any lone pairs.

So let's go ahead and do that.

So we're going to ignore the lone pair of electrons

on the sulfur when we're talking about the shape.

So if we ignore the lone pair and we actually

turn this molecule on its side-- so let's go ahead and do that.

We're going to put our sulfur here.

If we turn it on its side, the axial fluorines

would now be horizontal.

Now, it's horizontal like that.

So I'll go ahead and put in my-- so these

are the two fluorines that used to be axial there.

And my two fluorines that were equatorial,

they would look something like this.

And it helps if you actually build this molecule

with a MolyMod set.

So that would be what the molecule kind of looks

like here, and we call this a seesaw shape.

So this is a seeshaw shape or geometry.

And let's think about why.

So if you've ever been on a playground and used a seesaw--

I'm going to draw a little kid here

on one side of our seesaw like that.

And so if the little kid puts his weight

on this side, of course, this side of the seesaw salt

would go down.

And then, this side of the seesaw would go up.

So just a little bit of intuition

as to why you would call this a seesaw shape.

All right.

So I think we'll have to stop there.

In the next video, we'll do two more examples

of molecules and ions that have five electron clouds.