VSEPR Theory: Introduction

Summary

This video introduces the VSEPR theory, which helps predict the 3D molecular geometry from 2D Lewis structures. It explains that electron pairs repel each other, influencing molecular shape. Examples like BeCl2, CO2, BF3, SO2, CH4, NH3, and H2O illustrate various shapes: linear, trigonal planar, bent, and tetrahedral, affected by lone pairs. The video also guides viewers to further resources for practice and understanding of more complex molecules.

Takeaways

• 🔍 VSEPR (Valence Shell Electron Pair Repulsion) theory helps predict the 3D molecular geometry based on 2D Lewis structures.
• 🌐 The theory is based on the principle that electron pairs repel each other and arrange to be as far apart as possible.
• 🏔 In a molecule like BeCl2, the central beryllium atom is surrounded by two chlorine atoms, resulting in a linear shape with a 180° bond angle.
• 🔑 The VSEPR theory states that molecules with two electron pairs around a central atom tend to form linear structures.
• 🔄 Regardless of whether the bonds are single, double, or triple, the linear shape with 180° angles is maintained.
• 📐 For molecules like BF3 with three electron pairs, the 3D shape is trigonal planar with bond angles of 120°.
• 🔬 If a molecule has three electron pairs and one lone pair, like in SO2, the shape is bent or V-shaped with angles less than 120°.
• 🛑 The presence of lone pairs affects the molecular shape, causing bond angles to be smaller than those in molecules with only bonding pairs.
• 🔄 CH4, with four bonding pairs, forms a tetrahedral shape with bond angles of 109.5°.
• 🔄 NH3, with three bonding pairs and one lone pair, forms a trigonal pyramidal shape with bond angles slightly less than 109.5°.
• 💧 H2O, with two bonding pairs and two lone pairs, has a bent shape with bond angles of approximately 105°.

Q & A

What is VSEPR theory?

VSEPR theory stands for Valence Shell Electron Pair Repulsion theory, which is a model used to predict the geometry of individual molecules based on the repulsion between electron pairs in the valence shell of the central atom.

Why do electron pairs repel each other?

Electron pairs repel each other because they both carry a negative charge, and like charges repel each other according to the principles of electrostatics.

What is the significance of the term 'linear' in the context of molecular geometry?

In molecular geometry, 'linear' refers to a molecular shape where all atoms are aligned in a straight line, with bond angles of 180° between the electron pairs.

How does the presence of a lone electron pair affect the molecular shape?

The presence of a lone electron pair affects the molecular shape by pushing the bonded atoms closer together, resulting in bond angles that are smaller than those in a molecule with only bonding electron pairs.

What is the difference between a trigonal planar and a bent molecular shape?

A trigonal planar shape has three atoms around the central atom with bond angles of 120°, while a bent shape has one lone electron pair and two bonded atoms, resulting in bond angles of approximately 116°.

What is the bond angle in a tetrahedral molecular geometry?

In a tetrahedral molecular geometry, the bond angle between any two adjacent bonds is 109.5°.

How does the number of electron pairs surrounding the central atom influence the molecular shape?

The number of electron pairs surrounding the central atom determines the molecular shape according to VSEPR theory, with two electron pairs leading to a linear shape, three to a trigonal planar or bent shape, and four to a tetrahedral or trigonal pyramidal shape.

What is the term for a molecule with four electron pairs around a central atom, three of which are bonding pairs?

A molecule with four electron pairs around a central atom, three of which are bonding pairs, is said to have a trigonal pyramidal shape.

Why does the molecule BF3 have a trigonal planar shape?

BF3 has a trigonal planar shape because the three bonding electron pairs around the central boron atom arrange themselves to be as far apart as possible, resulting in bond angles of 120°.

How does the VSEPR theory explain the molecular geometry of CO2?

According to VSEPR theory, CO2 has a linear molecular geometry because the two double bonds are considered as two electron pairs, which arrange themselves in a straight line with 180° bond angles.

Outlines

00:00

🔬 Introduction to VSEPR Theory

This paragraph introduces VSEPR (Valence Shell Electron Pair Repulsion) theory, which is a method for predicting the three-dimensional shapes of molecules based on their two-dimensional Lewis structures. The theory is based on the principle that electron pairs around a central atom will arrange themselves to be as far apart as possible, minimizing repulsion. The paragraph uses the example of beryllium chloride (BeCl2) to illustrate how the electron pairs in single bonds repel each other, leading to a linear molecular shape with a 180° bond angle. The concept is further explained by discussing how double and triple bonds are treated as single electron pairs for the purpose of determining molecular geometry.

05:01

📐 Trigonal Planar and Bent Molecular Shapes

The second paragraph delves into molecules with three electron pairs surrounding a central atom, such as BF3, which adopts a trigonal planar shape with bond angles of 120°. The paragraph explains that this shape arises from the electron pairs' desire to be as far apart as possible. It then contrasts this with molecules like SO2, which has two bonding pairs and one lone pair of electrons, resulting in a bent or V-shaped molecular geometry with bond angles of approximately 116°. The influence of lone pairs on molecular shape is emphasized, as they exert a greater repulsive force than bonding pairs, thus affecting the bond angles.

10:01

🌐 Tetrahedral and Trigonal Pyramidal Geometries

This paragraph discusses molecules with four electron pairs around a central atom, exemplified by methane (CH4), which has a tetrahedral shape with bond angles of 109.5°. The tetrahedral geometry is explained as the arrangement that maximizes the distance between four electron pairs. The paragraph then explores how the presence of a lone pair, as in ammonia (NH3), alters the geometry to a trigonal pyramidal shape with bond angles slightly less than 109.5°. The concept is further illustrated by comparing the trigonal pyramidal shape with the tetrahedral one, highlighting the impact of lone pairs on the molecular geometry.

15:03

💧 Bent Molecular Shape of Water

The fourth paragraph focuses on water (H2O), a molecule with four electron pairs around the central oxygen atom, including two lone pairs. It explains how the presence of two lone pairs results in a bent molecular shape with bond angles of approximately 105°. The paragraph contrasts this bent shape with the trigonal pyramidal shape that arises when there is one lone pair and three bonding pairs. The summary emphasizes the cumulative effect of lone pairs on bond angles, with more lone pairs leading to smaller angles due to their greater repulsive influence.

20:04

🎓 Conclusion and Further Learning with VSEPR

The final paragraph concludes the introduction to VSEPR theory and suggests next steps for further learning. It recommends watching a practice problems video to apply the theory to various Lewis structures and predicts three-dimensional shapes. The paragraph also mentions common mistakes to avoid and additional videos covering more complex molecules with five or six electron pairs, such as trigonal bipyramidal and octahedral geometries. The summary encourages a structured approach to mastering VSEPR theory and its applications.

Keywords

💡VSEPR Theory

VSEPR stands for Valence Shell Electron Pair Repulsion theory. It is a model used in chemistry to predict the shapes of molecules by minimizing electron pair repulsions. In the video, VSEPR theory is central to explaining how molecules arrange themselves in three dimensions. The theory is used to describe the 3D shapes of molecules like BeCl2, CO2, and H2O, where electron pairs repel each other to achieve the lowest energy state.

💡Lewis Structures

Lewis structures are diagrams that represent the valence electrons of atoms within a molecule. They are used to determine the bonding between atoms and the arrangement of non-bonding electron pairs. The video script uses Lewis structures as a starting point to apply VSEPR theory and predict the 3D shapes of molecules, such as BF3 and SO2.

💡Central Atom

The central atom in a molecule is the atom that is bonded to the most other atoms or has the most non-bonding electrons. In the context of the video, the central atom is crucial for determining the molecular geometry. For instance, in BeCl2, the beryllium atom is the central atom, and its arrangement with chlorine atoms forms a linear shape.

💡Electron Pair Repulsion

Electron pair repulsion is the principle that electron pairs in the valence shell of an atom will arrange themselves to be as far apart as possible to minimize repulsion. This concept is fundamental to VSEPR theory and is used in the video to explain why molecules adopt certain shapes, such as the linear shape of CO2 where the electron pairs are 180° apart.

💡Linear Molecule

A linear molecule is one where the atoms are arranged in a straight line. The video explains that linear molecules form when there are two electron pairs (bonds) around a central atom, pushing each other to be as far apart as possible, resulting in a 180° bond angle, as seen in BeCl2.

💡Trigonal Planar

Trigonal planar is a molecular shape where the central atom is surrounded by three electron pairs, which are arranged in a flat, triangular shape. The video uses BF3 as an example, where the boron atom is surrounded by three fluorine atoms, each 120° apart, forming a trigonal planar shape.

💡Tetrahedral

Tetrahedral is a molecular geometry where the central atom is surrounded by four electron pairs, which are arranged to be as far apart as possible, forming a three-dimensional shape with bond angles of 109.5°. The video describes methane (CH4) as having a tetrahedral shape due to the four C-H bonds.

💡Trigonal Pyramidal

Trigonal pyramidal is a shape where a central atom is surrounded by three bonding electron pairs and one lone pair of electrons. The video explains that this shape is similar to tetrahedral but with one of the vertices replaced by a lone pair, resulting in bond angles slightly less than 109.5°, as seen in NH3.

💡Bent Molecular Shape

A bent molecular shape, also known as V-shaped, occurs when there are four electron pairs around a central atom, with two of them being lone pairs. The video describes water (H2O) as having a bent shape due to the two lone pairs on the oxygen atom, which push the hydrogen atoms closer together, resulting in a bond angle of approximately 105°.

💡Octet Rule

The octet rule states that atoms tend to form enough bonds so that they have eight electrons in their valence shell, giving them the electron configuration of a noble gas. The video mentions that some atoms, like beryllium in BeCl2, are exceptions to the octet rule, as they can be stable with fewer than eight electrons in their valence shell.

Highlights

Introduction to VSEPR theory as a tool to predict 3D molecular shapes from 2D Lewis structures.

Explanation of the concept that real molecules have 3D structures that are more complex than 2D representations.

VSEPR stands for Valence Shell Electron Pair Repulsion, emphasizing that electron pairs repel each other and influence molecular shape.

Electrons or pairs of electrons tend to stay as far apart as possible due to their negative charges.

Linear molecular shapes occur when two atoms bond to a central atom, with bond angles of 180°, as seen in BeCl2 and CO2.

The shape of a molecule is influenced by the repulsion between electron pairs in the bonds.

Double and triple bonds are treated the same as single bonds when determining molecular shape.

Trigonal planar shape forms when three atoms surround a central atom, with bond angles of 120°, as demonstrated with BF3.

SO2 has a bent shape due to lone pairs on the central atom pushing bonded atoms closer together.

Molecules like CH4, with four bonds around a central atom, form a tetrahedral shape with bond angles of 109.5°.

Trigonal pyramidal shape appears when there are three bonds and one lone electron pair, such as in NH3, with bond angles of about 107°.

Water (H2O) forms a bent shape due to two lone electron pairs on the oxygen atom, reducing bond angles to around 105°.

Lone pairs exert more repulsion than bonding atoms, reducing bond angles in structures like NH3 and H2O.

Differences in bent molecular shapes arise from having either three or four things around a central atom, with lone pairs playing a key role.

Follow-up videos include VSEPR practice problems, common mistakes, and advanced molecular shapes involving five or six surrounding atoms.

Transcripts

this video is an introduction to Vesper

Theory Vesper theory is a set of rules

that we can use to look at a

two-dimensional leis structure of a

molecule and figure out what the

molecule would look like in three

dimensions like this cuz molecules are

like actually things in real life so

they'd have three-dimensional structures

that are often more complex than we can

draw

in two Dimensions here so let's start

take a look at some Lew structures and

figure out what the 3D shapes of those

molecules would be here's our first

example burum D chloride we got a burum

atom it's a central atom surrounded by

chlorin on either side note that burum

here is an exception to the Octan rule

which means that it's happy to have

fewer than eight electrons in its veence

shell when burum is making two bonds

like it does here with chlorine it has

only four electrons in its veence shell

and it's perfectly happy with that so

just keep it in mind but it doesn't have

any important bearing on what the vasper

shape

is so a little bit about the Vesper

rules here Vesper stands for veence

Shell electron pair repulsion which is a

really fancy way of saying that

electrons or pairs of electrons want to

push away away from each other and want

to be as far away as possible from each

other and that kind of makes sense cuz

electrons have negative charges so

opposite charges repel and obviously

these things are going to want to be far

away from each other let's look at what

bearing this has on the

three-dimensional shape of a molecule so

in buril here where are these veent

shell electrons that want to push away

from each other well the veence shell

electrons are in these bonds I've often

said that you can think about coal bonds

as if they're hands from the atoms with

electron Pairs and that these hands are

connected because they're both holding

on they're both sharing this pair of

electrons so we could draw burum D

chloride like this where we have a hand

from the burum a hand from the chlorine

coming together to hold to share this

pair of electrons so this just

reinforces the idea that there is a pair

of electrons in each one of these bonds

that shared between the atoms okay so as

we said from Vesper these electrons want

to be as far away from each other as

possible they want to

repel so how is this going to influence

the 3D Shape of this molecule how can

these B bonds arrange each other in

arrange themselves in 3D so that they

are as far away from each other as

possible the three-dimensional shape of

burum D chloride is going to look like

this we've got a burum here in the

middle and then we have these two

chlorians on either side and all three

atoms form a line they're all in a row

here a straight

row we call this a linear molecule which

means line so that kind of makes sense

and let's look at the angles

here the angles between these two bonds

going to be

180° so

180° between these two bonds is how the

electrons that are in these two bonds

it's how they can be as far away from

each other as possible so we can say

that this linear shape that we have here

is the way that two things are going to

surround elv around a central atom

Central atom is here then we've got

these two things that are two bonds and

the two bonds are as far from each other

as possible in this linear

shape now in burum D chloride I'm

talking about single bonds here but it

actually doesn't matter whether we got

double bonds or triple bonds for example

CO2 has a shape like this where there's

a double bond here and a double bond

here but there are electrons in both of

these bonds and so each one of these

double bonds they just count as a bond

so for CO2 I still consider it as just

two things around a central atom so CO2

it's going to have this linear shape

here too this will be the carbon and

these will be the two oxygens they'll be

180° apart so double bonds don't worry

it's just two things one to around a

central atom okay just to drive this

point home trip bonds it's the same

thing got a triple bond here a single

Bond here I just consider this to be two

things around a central atom one 2 so

hcn is going to have this linear shape

as well with these two Ang these two

bonds being

180° apart so we always get a linear

molecule 180° whenever we have just two

things surrounding a central atom now

let's take take a look at some molecules

where we have three things that are

surrounding a central atom here in BF3 I

have a central atom surrounded by three

bonds to other atoms and in this case

Boron like buril before is an exception

to the octat rule here Boron when it's

making three bonds has six veence

electrons it's totally happy with that

so as I said earlier when we were

talking about burum is that we can think

about these bonds between the Boron and

the Florine here as hands that are

sharing an electron pair and the

electron pairs in each one of these

bonds push against each other and they

want to be far away so when we have

three things the electron pairs in these

three bonds how do we arrange these so

that they are as far away from each

other as possible in

3D the molecule is going to look like

this I'm going to have these three atoms

1 2 3 surrounding a central atom and I'm

going to get the shape called trigonal

planer the trigonal comes from the fact

that there are three one two three

things that's just what trigonal means

and planer because take a look at this

these atoms are all arranged in this

plane all right they're all in a

straight plane here now what are the

angles between the atoms in a trigonal

planer shape they are all

120° so the angle between here and here

is 120 here and here and here and here

so that is what BF3 would look like in

three dimensions now just as before it

doesn't matter whether we're talking

about double bonds single bonds triple

bonds it's all the same okay so in ch2o

here I have three things surround

rounding the central atom I got a double

bond a single Bond and a single Bond but

it's still just three things that want

to be as far from each other as possible

so that means that ch2o here is going to

have the same shape in three dimensions

as BF3 does it's going to be a trigonal

planer molecule with

180° between each pair of

bonds now next thing we're going to do

is we're going to look at some molecules

that have three things around them okay

but these three things are not all

bonds here's an example of this

S2 okay it's got three things around

this Central atom it's got a bond here

that's one thing a double bond here

that's two things but then it's got this

unshared electron pair up here these

three things all have electrons in them

so they all want to push away from each

other

so what shape is SO2 going to have in

three

dimensions if you think it's linear in

the three of these atoms are all lined

up in a row that's not right because

you're not taking this unshared electron

pair into

account to figure out the shape of this

let's go back to this trigonal planer

molecule okay in this trigonal planer

molecule we had three things around a

central atom it's just they were all

other atoms okay so this is how you

arrange three things around a central

atom to be far away from each other now

in SO2 we're going to get a shape that's

very similar except it's that one of

these atoms from the trigonal planer

shape is going to have been replaced by

an unshared electron pair but otherwise

look at how similar they are okay it's

atom atom atom atom it's just this atom

here has been replaced by this unshared

electron pair but these at atoms are

still in the same place because this

unshared electron pair pushes the atoms

away from each other just like this atom

did okay so they're based on the same

shape three things around a central atom

it's just one of these from the trigonal

planer shape has been replaced by an

unshared electron pair okay so this

molecule here we call this a bent

molecule because instead of being in a

straight line the atoms are arranged in

this kind of bent shape looks like

someone just grabbed it and bent it like

that so what are the angles going to

look like in the bent molecule well in

the trigonal planer molecule over here

when you had three atoms the angles

between any two Bonds were

120° in the bent molecule though it

turns out that this unshared electron

pair here pushes harder against these

two atoms than the atom up here would

okay and so that means that the angle

between these two bonds is going to be a

little bit less than 120 because the

atoms are getting pushed closer together

it's going to be less than 120 it's

going to be more like about

116° between the two of these just once

again because this unshared electron

pair is pushing harder than this atom so

instead of 120 they're pushed harder

pushed closer together and it's more

like

116 but here's a point we get the bent

molecule when one of these three atoms

from the trigonal planer is replaced by

a lone electron pair so always keep your

eye on these lone electron pairs because

they have a very significant impact on

the shape that a molecule is going to

end up having now let's move on to some

molecules that have four things around

the central atom CH4 here has four

things around a central atom and they

are all bonds to other atoms so each of

these bonds contain a pair of shared

electrons and that means that the bonds

all want to push against each other and

be as far from each other as

possible this molecule CH4 is going to

have this shape in three dimensions okay

this is called a tetrahedral shape and

it's how you arrange four bonds as far

away from each other in 3D as possible

okay tetrahedral

and in the tetrahedral molecule there

are

109.5° between any two bonds that are

next to each other in this molecule so

109.5 here

109.5 and so on so four things four

things around a central atom you get a

tetrahedral shape with 109.5 degrees

between Each Bond now NH3 here also has

four things around a central atom but

not all of them are bonds to other atoms

okay so we have 1 2 three bonds and then

a fourth thing that's alone electron

pair so what's its shape going to look

like in three dimensions I'm going to go

back to this tetrahedral shape for just

a minute because this is how we arrange

four things around a central atom when

they're all other atoms okay but in this

shape they're not all other atoms

okay so NH3 is going to end up having

this shape which is called a trigonal

pyramidal shape look at how similar it

is to the tetrahedral shape okay I'm

sort of showing them on their sides here

it's just that the atom that was up here

when we had four atoms around the

central atom has been replaced by an

unshared electron pair here okay so

we've got three atoms three atoms are

the same between this and this and it's

just this atom has been replaced by an

unshared electron pair so this has a

shape that we call trigonal pyramidal

and we call it trigonal pyramidal

because if you look at it from its side

it kind of looks like a pyramid okay got

these three atoms pointing down all

right now for angles in I don't know

quite I would to put this I put it up

here I guess for atoms in the trigonal

pyramidal for angles in the trigonal

pyramidal mod

molecule we'll remember that in the

tetrahedral we have

109.5° between all of the bonds but a

trigonal pyramidal just like we saw with

a bent molecule the unshared electron

pair here pushes a little harder against

these two bonds than an atom would and

so that means that the angle between

these bonds is pushed a little tighter

and so it's smaller than

109.5 for a trigonal pyramidal molecule

like

NH3 the bond angle is more like

107° little less than

109.5° so four things if you have four

things around a central atom but three

of them are bonds and one of them is a

lone electron pair you end up with a

shape it's called trigonal pidal that

looks like this okay one more example

and then we're done with done with a

Vesper video here's the last molecule

we're going to look at water H2O okay

this thing has four things around a

central atom two of them are bonds one

two and two of them are lone electron

pairs one two so what's its 3D shape

going to look like how can we arrange

these four things as far away from each

other as possible in three dimensions as

before I'm going to look back at my tetr

molecule which shows how I arrange four

atoms or four bonds as far away from

each other as possible in

3D for H2O though only two of the four

things are bonds the other two are lone

electron pairs okay so that means that

I'm going to end up with a

shape like this okay I've got my two

atoms down here hydrogen and hydrogen

but then I've got my two lone electron

pairs up

here look at how this is similar to the

tetrahedral molecule if I look at them

from the side okay it's got I've got

atom atom and atom and atom okay it's

just these two atoms from the

tetrahedral molecule have been replaced

by these two lone electron pairs from

the water molecule these unshared

electron pairs on the oxygen here okay

so look at that from the top

how they have very similar structures

it's just these two are missing and

they've been replaced by the lone

electron

pairs we say that this molecule has a

bent shape because these Mo these atoms

here are not in a straight line but

they're bent like this now what what are

the angles here let's look at the

tetrahedral again which had

109.5 then in the trigonal pyramidal

when we had one electron pair it pushed

the bonds a little bit closer together

so we had about 107° between them a

little less than 109.5 now when we have

bent instead of one unshared electron

pair like in the trigonal planer I have

two lone electron Pairs and so the

combination of those two is going to

push the atoms even a little bit closer

so in a bent molecule like water the

angle between them is going to be 105

degrees about 105 degrees Which is less

than 107 Which is less than 109.5 so the

more unshared electron pairs you add the

tighter the two or more atoms get pushed

together okay so if you got two bonds

and two lone electron pairs around a

central atom you're going to have this

bent shape now there's just one thing

that I want to say about this bent shape

okay there are two ways that we can get

a molecule with a Ben shape but they're

different okay we can get a bench shape

when we have three things around a

central atom and one of them is an

unshared electron pair then we get

something like SO2 where we have

something that's a little less than 120°

between these another way to get a bench

shape is when we have four things around

a central atom but two of them are lone

electron Pairs and in that case because

we have four things everything's a

little tighter we have an angle of 105

degrees that's a little less than

109.5 so you just finished watching this

Vesper video where should you go from

here well the first thing that it's

important to know is that there's a

difference between watching the video

and actually being able to look at a

leis structure and figure out what the

3D Shape of that molecule would be so

the first video that you should watch is

this Vesper practice problems video

where we'll go through a whole bunch of

Lewis structures and go through the

steps to figure out what the

three-dimensional shapes will be so

that's definitely the next thing you

should watch now there are some common

mistakes that students often make when

they're learning Vesper so I made a

video on that called Vesper common

mistakes watch that after you've done

the practice problems to make sure that

you're not falling into any of the

common traps that tend to trip people up

when they're learning Vesper now maybe

this is all of the stuff that you have

to learn for Vesper and this is

everything up to molecules that have

four things around a central atom but

depending on what you have to know you

might have to know molecules where there

are five things like this around the

central atom or where there are six

things like this around a central atom

so I made some other videos on this sort

of stuff okay I made the video a video

on the trigonal bipyramidal family which

are all of the molecules that have five

things around a central atom and then

there's another video on the octahedral

family which are the molecules that all

have six things around a central atom

and now finally after you've watched the

video on the trigonal bipyramidal and

the octahedral you can do the Vesper

practice problems uh for these Advanced

structures where you where you'll go

over the uh the molecule shapes I talk

about in this video in this video so

this uh this might look a lot this might

look like a lot but if you go through it

it should really give you a solid

foundation with this threedimensional V

asper stuff