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Oxygen (O2) is a diatomic molecule, meaning it consists of two oxygen atoms bonded together. To determine if a molecule is polar or nonpolar, we need to consider two main factors: the electronegativity of the atoms and the geometry of the molecule.
Oxygen (O2) is a diatomic molecule, meaning it consists of two oxygen atoms bonded together. To determine if a molecule is polar or nonpolar, we need to consider two main factors: the electronegativity of the atoms and the geometry of the molecule.
1. Electronegativity: This is a measure of how strongly an atom attracts electrons in a chemical bond. In the case of O2, both oxygen atoms have the same electronegativity since they are the same element. Therefore, there is no difference in electronegativity between the two oxygen atoms.
2. Molecular Geometry: The shape of a molecule can also determine its polarity. For a molecule to be polar, it must have an asymmetric shape that allows for an uneven distribution of charge. However, in the case of O2, the molecule is linear, with the two oxygen atoms sharing electrons equally in a double bond.
The double bond between the oxygen atoms is described by the following Lewis structure:
Since the two oxygen atoms have the same electronegativity and the molecule is symmetrical, the electrons are shared equally, and there is no permanent dipole moment. A dipole moment occurs when there is a separation of charge within a molecule, leading to a molecule having a positive end and a negative end.
In summary, because O2 has a linear geometry and the atoms involved have equal electronegativity, there is no net dipole moment. Therefore, O2 is a nonpolar molecule.
1. Electronegativity: This is a measure of how strongly an atom attracts electrons in a chemical bond. In the case of O2, both oxygen atoms have the same electronegativity since they are the same element. Therefore, there is no difference in electronegativity between the two oxygen atoms.
2. Molecular Geometry: The shape of a molecule can also determine its polarity. For a molecule to be polar, it must have an asymmetric shape that allows for an uneven distribution of charge. However, in the case of O2, the molecule is linear, with the two oxygen atoms sharing electrons equally in a double bond.
The double bond between the oxygen atoms is described by the following Lewis structure:
Since the two oxygen atoms have the same electronegativity and the molecule is symmetrical, the electrons are shared equally, and there is no permanent dipole moment. A dipole moment occurs when there is a separation of charge within a molecule, leading to a molecule having a positive end and a negative end.
In summary, because O2 has a linear geometry and the atoms involved have equal electronegativity, there is no net dipole moment. Therefore, O2 is a nonpolar molecule.
Dipole Moments of Molecules
Exercise 4 Part c - Predicting the polarity of molecules (CCl₂F₂)
Exercise 3 - Is SCl₂ a polar or nonpolar molecule?
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In a previous video,
we talked about the dipole moments of individual polar covalent bonds.
In this video, we will determine the dipole moments of molecules.
A polyatomic molecule is polar,
that means it has a dipole moment if the sum of
the dipole moments of the individual bonds is non-zero.
In order to find out whether the sum is 0 or non-zero,
we need to know the geometry of the molecule.
Now we're going to deal with a number of examples to illustrate this point.
Let's start off with CO_2.
CO_2 belongs to the AX_2 category of VSEPR.
Now the CO bond itself is a polar covalent bond.
The C is slightly positive and the O is slightly negative.
So the dipole moment points towards O.
Why is O more negative than C?
That's because the electronegativity of
O is greater than the electronegativity of C. C is 2.5 and O is 3.5,
so the dipole moment of the CO bond points towards O.
Now let's look at the CO_2 molecule.
It's linear.
The C is AX_2,
so the sum of the dipole moments is 0.
We have 1 going this way and 1 going that way, and they cancel.
So CO_2 is a non-polar molecule.
Now let's look at water, H_2O,
which is AX_2E_2 in VSEPR language.
Now the HO bond itself is polar.
H is slightly positive,
O is slightly negative,
with a dipole moment pointing towards O.
Since O is more electronegative than H,
H is 2.1 and O, 3.5.
The dipole moment is pointing in this direction.
Now as I've said many times in previous videos,
water is a bent molecule with the angle HOH=104.5 degrees.
The sum of the dipole moments is non-zero. Let's draw it.
O, H, H,
1 dipole is pointing this direction,
1 in this direction,
and the sum of the dipoles is in this direction,
so H_2O is a polar molecule.
The third example is BF_3,
which belongs to AX_3 category of VSEPR.
The BF bond is a polar covalent bond with a dipole pointing
towards F. That's because F is more electronegative than B.
B has electronegativity of 2 and F of 4.
Now, what happens when we get to the molecule?
We've said before that BF_3 is a planar molecule with the FBF angle 120 degrees.
B is in the center, 3 Fs.
Now if we add the 2 dipoles pointing down the way,
we get a dipole of equal value that bisects the FBF angle.
Here is the resultant of these 2 vectors.
The top BF bond,
for it, the dipole is in the opposite direction.
When we sub the red vector and the black vector, we get 0.
That means BF_3 is a non-polar molecule.
The fourth example, COCL_2,
which again is AX_3.
But this time the CCL dipole moment is different from the CO dipole moment.
The molecule is polar.
We have CO that dipole in this direction
and CL resultant of these 2 dipoles.
The same lens and in the downwards direction,
but it's a different lens,
it has a different amplitude than the CO vector pointing towards the top.
The sum is non-zero.
That means that COCL_2 is a polar molecule.
It has a resultant dipole moment,
a net dipole moment.
The fifth example is CCL_4,
which belongs to the AX_4 category of VSEPR theory.
Now we'll see that the sum of the dipole moment is 0.
Each CCl bond has a dipole moment,
but the sum is 0.
The way to see that is to draw the molecule inside a cube.
We can draw a tetrahedral molecule inside a cube.
We have C in the center,
a bond pointing towards this vertex.
We take the sum of these two,
we get a vector pointing down.
Now we have another 2 bonds,
1 pointing to this vertex,
and 1 pointing to this vertex.
This makes our tetrahedron.
The sum of these two points upwards.
It's exactly equal and opposite to the 1 pointing down in the way,
so the sum is 0.
The sum of bond dipole moments is 0 and that means that CCL_4 is a non-polar molecule.
If 1 of the 4 X atoms is different from the other 3,
for example, CH_3CL,
then the molecule will be polar.
Now, why is that so?
CH has a very small dipole moment,
0.122 debye, whereas CCL has much greater dipole moment of 1.6 debye.
If we draw the molecule,
then the dipole moments of the CH bonds will be very small, pointing downwards,
and the much larger dipole of the CCL bond will be pointing upwards,
giving us a net dipole moment in the same direction as the CCL bond.
This molecule is polar.
Now if we follow the same reasoning,
it's easy to see that AX_5 and AX_6 molecules are non-polar provided of course,
that all of the X atoms are the same.
That means PCL_5, for example,
or SF_6 are non-polar molecules.
In this video, we talked about the dipole moments of molecules.
This video explains how to determine the dipole moments of molecules. It starts by discussing the dipole moments of individual polar covalent bonds and then moves on to the sum of the dipole moments of the individual bonds to determine whether the molecule is polar or non-polar. Examples are given to illustrate the point, such as CO_2, H_2O, BF_3, COCL_2, and CCL_4. It is concluded that if one of the X atoms is different from the other 3, the molecule will be polar, while AX_5 and AX_6 molecules are non-polar.
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