Molecular solids are composed of discrete molecules held together by intermolecular forces. Because these interactions are relatively weak, molecular solids tend to be soft and have low to moderate melting points. Molecular solids are also poor conductors of electricity because their valence electrons are tightly held within each individual molecule. Created by Sal Khan.
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- In the last part of the video (6:00onwards) Sal asks to compare the melting points of molecular Iodine (I2) to that of molecular Chlorine (Cl2)
The video states that because Iodine has a large atomic radius and is therefore a larger atom, it is more polarizable with greater dispersion forces and therefore has a higher melting point that the chlorine molecule....
But if we compare the atomic radii of Cl vs. I, because I is in a lower period (has more shells) than Cl, we can assume that Iodine's electrostatic force is lower than that of Cl (because it has a greater atomic radius that is inversely related to electrostatic force); thereby suggesting that Iodine should have a lower melting point than the molecular Chlorine...
I don't understand Sal's explanation of why Iodine's melting point is indeed higher than Chlorine's. How does it contrast to the logic/answer I deduced? How/why am I looking at it incorrectly?
Please note that - I haven't studied yet intermolecular forces - because I am following the Chemistry syllabus and no videos on these topics have been shown yet (already posted a tip suggesting something be done to address this as I am surely not the only one that was lost by the sequential order and content of this video).(8 votes)
- You note you haven’t got to intermolecular forces yet, it may be a good idea to wait until you get to those videos then maybe revisit this question.
You mentioned electrostatic force so Coulomb’s law but that’s relevant to ionic compounds, but this video is dealing with covalently bonded molecules.
When you get to intermolecular forces you’ll learn about dispersion forces which are stronger with more electrons and larger atoms. As iodine atoms are significantly larger and have significantly more electrons than Cl atoms, this is why I2 has a higher melting and boiling point than Cl2.(6 votes)
- Why is I more polar just because it is larger? I remember previous videos mentions the electronegativity of elements decreases as they go from top to bottom; thus, Cl should be more polar than I. Right?(2 votes)
- Electronegativity and polarizability are are different effects at play here. Electronegativity describes how well an atom attracts a pair of shared electrons in a covalent bond to itself. While polarizability describes essentially how well charges can be distorted on an atom in the presence of an external electric field. Electronegativity is used more when dealing with dipole-dipole intermolecular forces, while polarizability is used when dealing with London dispersion forces.
In a sample of molecular I2, every atom has the same electronegativity value because they are all the same element. Every molecule of I2 has no dipole moment because both atoms of iodine pull on the electrons in the covalent bond equally as much making each molecule non-polar. Same goes for a sample of molecular Cl2. So even though chlorine is more electronegative than iodine, it doesn't matter because both samples are nonpolar because they are composed of the same respective chemical elements.
Polarizability views the atoms as spheres with negative electric charges on the surface of the sphere from the electrons. Having a larger atomic radius in the case of iodine because of its higher shell number means the electrons are more easily able to be pushed around the atom by outside electric fields. These outside electric fields coming form the electrons of other I2 molecules. Since the electrons are more easily able to move to one side of an I2 molecule because of the repulsion from another I2's electrons this creates a temporary dipole moment in the molecule. This type of temporary dipole moment between non polar molecules (all molecules though really) is what we call London dispersion forces. A bigger atomic radii in iodine as compared to chlorine is what causes this increased polarizability, hence greater London dispersion force and a higher melting point. Hope that helps.(5 votes)
- Sal only uses London Dispersion forces in his examples. But it is possible to have molecular solids held together with other IMFs, right? For example, solid H2O (ice), which consists only of non-metals, is held together with hydrogen bonds...? Right?(2 votes)
- Why are molecular solids insoluble in water?(1 vote)
- Well that's not strictly true, there are examples of molecular solids which are soluble in water. The simplest being sucrose, otherwise known as table sugar, and caffeine which are found in beverages composed mostly of water. Really what determines whether a solid will dissolve in water or not is if it is polar.
Chemicals dissolve in other chemicals based on their polarity and generally follows the rule 'like dissolves like'. Where polar solids dissolve in polar liquids, nonpolar solids dissolve in nonpolar liquids, but nonpolar solids do not dissolve in polar liquids (also polar solids do not dissolve in nonpolar liquids). Water is polar molecule and will generally only allow other polar molecules to dissolve into it. So a molecular solid like solid iodine, I2, will not dissolve in water because iodine is nonpolar, but sucrose will dissolve in water because sucrose is sufficiently polar (it has a lot of polar -OH groups on it).
Hope that helps.(1 vote)
- How do you distinguish between covalent network solids and molecular solids when looking at the chemical representation? like SiC and SiO or something like that?(1 vote)
- Are molecular solids crystalline or amorphous? Why is that?(1 vote)
- Crystalline means there's a repeating pattern to your solid and forms a crystal lattice, while amorphous means there is no pattern and they do not form a defined geometric shape. Molecular solids would belong to the crystalline group which form lattices weaker than that of lattices in ionic solids. This is because ionic crystals are held together by strong electrostatic attractions while molecular solids are held together by comparatively weak intermolecular forces. Hope that helps.(1 vote)
- [Instructor] So let's talk a little bit about molecular solids. So just as a little bit of review, we've talked about ionic solids where ions form these lattices. So those might be the positive ions right over there, and then you have your negative ions. And the negative is attracted to the positive. The positive is attracted to the negative. And I'm just showing a two-dimensional version of it, but it forms a three-dimensional lattice. So that's an ionic solid. We have also seen metallic solid where you have metals that all contribute some valence electrons to the sea of electrons. So what you end up having is essentially these positive cations that are in this sea of electrons. And we've talked about those properties, very good at conducting electricity, malleable, et cetera. Now, what we're gonna do is talk about what happens when you have nonmetals. So the nonmetals, you can see in yellow right over here, also includes hydrogen. Now, of course, noble gases are also nonmetals, but they're not reactives. So we're gonna talk about the reactive nonmetals. They can form molecules with each other. For example, one iodine can bond to another iodine with covalent bonds. So you could have a molecule like I2. You have things like carbon dioxide. Each carbon can bond to two oxygens. These are each molecules formed due to covalent bonds between nonmetals. Now, when we talk about molecular solids, we're talking about putting a bunch of these together. So let's say putting a bunch of iodine molecules together, and the intermolecular forces at a sufficiently low temperature are sufficient to hold together those molecules as a solid. So what do I mean by that? Let's look at a few examples. This right over here is a picture of solid iodine, and the way it's made up is you have these iodine molecules. Now, each of these iodine molecules are formed by a covalent bond between two iodine atoms. Now, the reason why it's a solid is there's enough dispersion forces. We talked about these London dispersion forces that are formed by temporary dipoles inducing dipoles in neighboring molecules. For example, just by random chance, for a moment, you might have more electrons on this end of this iodine molecule, creating a partially negative charge. And then that means some of the electrons on this end of this neighboring iodine molecule might be repulsed by that negative charge, so it forms a partially positive charge. And so you have a temporary dipole inducing a dipole in the neighboring molecule, and then they'll be attracted to each other, and we've talked about that as London dispersion forces. And at a sufficiently low temperature, that can keep them altogether as a solid. Now, it's important to point out, I keep saying sufficiently low temperature because these molecular solids, because they are only held together not by the covalent bonds, the covalent bonds hold together each of the molecules, but the molecules are held together by these fairly weak dispersion forces. They tend to have relatively low melting points. For example, solid iodine right over here has a melting point, has a melting point of 113.7 degrees Celsius. And I know what you're saying. That's not that low. That's higher than the temperature at which water boils. It would be quite uncomfortable for any of us to be experiencing 113.7 degrees Celsius. But this is relatively low when you talk about solids. Think about the temperatures it requires to melt, say, table salt. We've talked about that. Think about the temperatures it takes to melt iron. There, you're talking about hundreds of degrees, in certain solids, thousands of degrees Celsius. This is much lower. And so as a general principle, molecular solids tend to have relatively low melting points. Now, how good you think they're gonna be as conductors of electricity? Pause the video and think about that. Well, in order to be conductors of electricity, somehow charge needs to move through the solid. And unlike metallic solids, you don't have the sea of electrons that can just move around, so these tend to be bad conductors of electricity. If you wanna see another example of a molecular solid, this right over here is solid carbon dioxide, often known as dry ice. What you see here is each of these molecules, each carbon, is bonded to two oxygens. It has a double-bond with each of those oxygens. These are covalent bonds that form each of these molecules. But what keeps all of the molecules attracted to each other is, once again, those dispersion forces. And these forces between the molecules are so weak that solid carbon dioxide doesn't even really melt. It doesn't even go to a liquid state. If you heat it up enough to overcome these intermolecular forces, these dispersion forces, it will sublime, which means it goes directly from a solid to a gas state, and it does that at a very low temperature. It sublimes at negative 78.5 degrees Celsius. And if you've ever handled a dry ice, which I don't recommend you doing without gloves because it will hurt your skin if you do touch it, I actually did that recently at my son's birthday party, we were playing around with dry ice, you don't mess around with this thing because it is so incredibly cold. And at that temperature, it will go from a solid. It won't even melt to a liquid state. It will go straight to a gas state. Now, the last thing I wanna do is think about why different molecular solids will have different melting points. So let's compare, for example, molecular iodine to molecular chlorine. Each of these can form molecular solid. We looked at iodine a few minutes ago. Which of these would you think would form molecular solids with higher melting points? Pause the video and think about that. Well, as we talked about it, each of these molecules, they're formed by covalent bonds between two atoms, and what keeps the solid together are these dispersion forces. In an earlier videos, when we first talked about dispersion forces, we talked about temporary dipoles and induced dipoles, and they were likely to form between heavier atoms and molecules because they have larger electron clouds and are more polarizable. So if you compare molecular iodine to molecular chlorine, you can see that iodine is clearly made up of larger atoms and is therefore a larger molecule, which is more polarizable. So it's larger, which means it's more polarizable, generally speaking, polarizable, which means it has stronger, generally speaking, dispersion forces, stronger dispersion forces. Now, just as a reminder, these dispersion forces are between molecules. Each molecule has a covalent bond between two iodines, and then the dispersion forces are between the molecules. But because it has stronger dispersion forces, we would expect that a molecular solid formed by iodine is gonna have a higher melting point than a molecular solid formed by chlorine. And I actually do have the numbers here. So the melting point of a molecular solid formed by iodine, we've already talked about that, that's 113.7 degrees Celsius, while the melting point of a molecular solid formed by molecular chlorine has a melting point of negative 101.5 degrees Celsius, which is very cold, and so iodine has a higher melting point because of the stronger dispersion forces. Now, as I said, those dispersion forces are still not that strong. This is still not that high of a temperature compared to melting points of other types of solids we have looked at in the past.