Hello Chemistry 20! This is our seventh lesson in the chemical bonding unit. This is intermolecular forces.
In this lesson we will compare intermolecular forces. molecular forces with intra molecular forces we will explore three different types of intermolecular forces and work through three examples before starting the lesson it's important to remember fundamental principle in chemistry that similar charges repel each other and opposite charges attract Intra versus intermolecular forces. Intra-molecular forces refer to forces within a molecule. They are the forces of attraction that hold atoms together. They are, for example, covalent bonds.
As we can see in this picture, an intra-molecular force is the bond between two atoms For example, the bond between hydrogen and chlorine. Intermolecular forces occur outside of molecules. They are the forces of attraction between different molecules.
They are not bonds, they are just electrostatic forces of attraction. And as we can see in this picture, our intermolecular force is the force of attraction between two different molecules. They will attract each other and come close together.
The first type of intermolecular force is a dipole-dipole interaction. Remember, dipole is also known as a polar molecule. Therefore, this is a polar molecule-polar molecule interaction.
This force of attraction is between the slightly positive side of one polar molecule and the slightly negative side of another polar molecule. For example, if we look at the molecule HCl, there is a positive and there is a negative side to this molecule. Well, the negative side of this HCl molecule will be attracted to the positive side of another HCl molecule.
Remember, opposite charges attract. The negative side of this molecule is attracted to the positive side of this molecule. Two polar molecules attracted to each other, this is a dipole-dipole attraction.
Our next type of intermolecular force is a special type of dipole-dipole force. That is because it is the strongest intermolecular force. So again, this is still a polar molecule, polar molecule attraction. This is hydrogen bonding. This is when we have a force of attraction between a hydrogen atom bonded to an oxygen nitrogen, or fluorine.
And that hydrogen atom is attracted to the lone pair of electrons on another nitrogen, oxygen, or fluorine. If we look at our example picture, we can see that we have a hydrogen right here. It is bonded to an oxygen.
This hydrogen atom is attracted to the lone pair of electrons on another oxygen nitrogen, or fluorine. Therefore, this type of dipole-dipole force will qualify for hydrogen bonding, and this will be the strongest type of intermolecular force. Consider the following. Why is hydrogen bonding the strongest intermolecular force? Well, it's important that we take a look at how one of these molecules actually looks.
We know that the bond between oxygen and hydrogen is polar, and those electrons will be located much closer towards oxygen. Well, notice how the hydrogen is actually just an exposed proton. That is a very positive force.
And the lone pair of electrons on another oxygen is a very negative force. When we have a very positive and very negative force, we have a very strong force of attraction. That is why hydrogen bonding is the strongest intermolecular force.
And our last type of intermolecular force. London dispersion forces. When electrons are bonded they are constantly vibrating. The vibrating electrons can move out of their fixed location instantaneously for a brief moment in time causing a nonpolar molecule to become polar or a polar molecule to become even more polar.
This is the weakest type of intermolecular force and the reason why is because this is an instantaneous force of attraction the moment those electrons vibrate out of location they will just as fast vibrate back into their fixed location so we can see in this example that we have a cl2 molecule which is non-polar So it should not have any forces of attraction because there is no positive or negative side to this molecule. But if the electrons in this molecule vibrate and move out of their location instantaneously, this molecule can become slightly polar for a brief moment in time. The moment it becomes polar, it will have a slightly positive and slightly negative side, which can then be attracted to other molecules. So it's important to note that these types of forces, London dispersion forces, are instantaneous.
The moment an electron vibrates out of location and we have a force of attraction, those electrons will instantaneously vibrate back to their location and the force of attraction will be gone. This is an extremely weak force of attraction. It's important to note the larger the molecule, the more electrons the molecule has, which will result in more London dispersion forces.
Just to recap the strength of intermolecular forces. Our strongest type of intermolecular forces is hydrogen bonding. Then it's followed by dipole-dipole interactions.
And finally... weakest type of intermolecular force is London dispersion forces it's also important to note that intermolecular forces all of our intermolecular forces will be much weaker than intra molecular forces the bonds between atoms example number one determine the VSEPR diagram the polarity and the intermolecular forces present in NO2-. Pause the video and attempt this example.
First up, we need a proper Lewis dot structure. Remember, if you can't draw the Lewis dot structure, you cannot do anything moving forward. So as we can see, we have nitrogen as our center atom, and we have three electron groups around nitrogen.
A single bond, a double bond, and the lone pair of electrons. This is an ion, so we do show the brackets and the minus charge. So our VSEPR diagram will look like the following.
We have the lone pair up top and the two bonds going slightly downwards. We can then conclude that this molecule has a vent molecular shape, and this is going to be a polar molecule. It is asymmetrical. Therefore, if it's a polar molecule, it will experience London dispersion forces. and dipole-dipole intermolecular forces.
It will not qualify for hydrogen bonding because we do not have a hydrogen bonded to a nitrogen, oxygen, or fluorine. Example number two. Determine the VSEPR diagram, the polarity, and the intermolecular forces present in HNO.
Pause the video and attempt this example. First up is our VSEPR diagram. So we've skipped right past the Lewis dot structure, which again is the most important. But what we can see is nitrogen is our center atom, and we have a double bond to oxygen, a single bond to hydrogen, and a lone pair of electrons. So we can see that there are three electron groups around nitrogen, making this bent molecular shape.
the lone pair of electrons causes it to be bent and this molecule will be asymmetrical therefore it is a polar molecule and this molecule will have all three types of intermolecular forces all molecules have london dispersion forces this is a polar molecule so we'll have dipole dipole interactions, and we do have a hydrogen. atom bonded to a nitrogen, oxygen, or fluorine. So this molecule will qualify for hydrogen bonding. And finally, our last example, determine the VSEPR diagram, the polarity, and the intermolecular forces present in CH4.
Pause the video and attempt this example. Okay, first up is our VSEPR diagram. So we can see that carbon is our center atom, and we have four hydrogens bonded around the carbon. So there's four electron groups on the carbon, making this a tetrahedral molecular shape.
There are no lone pairs, so its molecular and geometrical shape will be tetrahedral. This molecule is symmetrical because we have all of the same atoms bonded around carbon. Therefore, this will be a nonpolar molecule.
And if it's a nonpolar molecule, it only has London dispersion forces. It is not polar, so therefore it cannot have dipole-dipole interactions. Consider the following.
Non-polar molecules only experience London dispersion forces. Polar molecules experience London dispersion forces and dipole-dipole interactions. While some polar molecules can experience London dispersion forces, dipole-dipole interactions, and hydrogen bonding.
But they must have a hydrogen atom. bonded to an oxygen, nitrogen, or fluorine. Moving forward, we will explore the physical properties of ionic and covalent compounds.