VSEPR Theory: Introduction
Summary
TLDRThis 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
- π VSEPR (Valence Shell Electron Pair Repulsion) theory helps predict the 3D molecular geometry based on 2D Lewis structures.
- π The theory is based on the principle that electron pairs repel each other and arrange to be as far apart as possible.
- π In a molecule like BeCl2, the central beryllium atom is surrounded by two chlorine atoms, resulting in a linear shape with a 180Β° bond angle.
- π The VSEPR theory states that molecules with two electron pairs around a central atom tend to form linear structures.
- π Regardless of whether the bonds are single, double, or triple, the linear shape with 180Β° angles is maintained.
- π For molecules like BF3 with three electron pairs, the 3D shape is trigonal planar with bond angles of 120Β°.
- π¬ 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 120Β°.
- π The presence of lone pairs affects the molecular shape, causing bond angles to be smaller than those in molecules with only bonding pairs.
- π CH4, with four bonding pairs, forms a tetrahedral shape with bond angles of 109.5Β°.
- π NH3, with three bonding pairs and one lone pair, forms a trigonal pyramidal shape with bond angles slightly less than 109.5Β°.
- π§ H2O, with two bonding pairs and two lone pairs, has a bent shape with bond angles of approximately 105Β°.
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 180Β° 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 120Β°, while a bent shape has one lone electron pair and two bonded atoms, resulting in bond angles of approximately 116Β°.
What is the bond angle in a tetrahedral molecular geometry?
-In a tetrahedral molecular geometry, the bond angle between any two adjacent bonds is 109.5Β°.
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 120Β°.
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 180Β° bond angles.
Outlines
π¬ 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 180Β° 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.
π 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 120Β°. 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 116Β°. 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.
π 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 109.5Β°. 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 109.5Β°. 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.
π§ 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 105Β°. 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.
π 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.
Mindmap
Keywords
π‘VSEPR Theory
π‘Lewis Structures
π‘Central Atom
π‘Electron Pair Repulsion
π‘Linear Molecule
π‘Trigonal Planar
π‘Tetrahedral
π‘Trigonal Pyramidal
π‘Bent Molecular Shape
π‘Octet Rule
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 180Β°, 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 120Β°, 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 109.5Β°.
Trigonal pyramidal shape appears when there are three bonds and one lone electron pair, such as in NH3, with bond angles of about 107Β°.
Water (H2O) forms a bent shape due to two lone electron pairs on the oxygen atom, reducing bond angles to around 105Β°.
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
180Β° so
180Β° 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
180Β° 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
180Β° apart so we always get a linear
molecule 180Β° 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
120Β° 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
180Β° 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
120Β° 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
116Β° 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
109.5Β° 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
109.5Β° 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
107Β° little less than
109.5Β° 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 107Β° 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 120Β°
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
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