Covalent network solids are composed of atoms covalently bonded together into a three-dimensional network or layers of two-dimensional networks. Due to the strength of the covalent bonds, covalent network solids have high melting points. Three-dimensional network solids (such as diamond or silica) are hard and rigid, whereas two-dimensional network solids (such as graphite) are soft due to the ease with which the network layers can slide past each other. Created by Sal Khan.
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- Are there any patterns in covalent network solids as opposed to molecular solids (e.g. Electronegativity difference)?. A lot of the seem to have elements in the region around the metalloids. Is there a reason for this?(3 votes)
- Well whether the solid is a network solid or molecular solid, they use covalent bonding. Nonmetals in the upper right corner of the periodic table engage in this type of bonding as opposed to the metals in the rest of the periodic table which engage in metallic bonding. This is just due to nonmetals having higher electronegativities than metals.
The main difference between the solids is that network solids are held together due to covalent bonding, while molecular solids are held together due to Van Der Waals forces like London dispersion forces. But as to why an element would rather form a network solid as opposed to a molecular solid, or vice versa, there is no clear pattern why. Atoms of the same element can form either types of solids often.
Take carbon for example, it can form diamond which is a network solid of carbons covalently bonded to each other forming a tetrahedral lattice. But then it can also form a variety of organic molecular solids like naphthalene.
Oxygen forms a molecular solid as a diatomic oxygen, but pair it with silicon and it makes the silicon dioxide lattice which is a network solid. Phosphorus can be a molecular solid as white phosphorus, but can also form a network solid called red phosphorus.
Hope that helps.(7 votes)
- How do I know if any specific molecule will form a molecular solid or a covalent network? For instance, how would I know that SiO2(s) will form a covalent network without already knowing; would I just have to memorize them?(1 vote)
- Yeah you'd have to memorize them. Sometimes we try to differentiate the formula of network solids by writing something like (SiO2)n or (C)n. The 'n' stands for many multitudes of these repeating units since a network solid is a polymer of these formula unit monomers.
Hope that helps.(9 votes)
- How do you tell whether a molecule is a covalent network solid or a molecular solid??(1 vote)
- Well firstly we don’t categorize molecules as either network or molecular solids; rather we describe solids using those terms.
A network solid is a solid where all the atoms are covalently bonded in a continuous network. Similar to an ionic solid where it’s a repeating pattern of formula units held together by ionic bonding, here they are held together by covalent bonds. This is also why we can’t describe a solid like a network solid as a molecule since there are no discrete separate molecules. It’s all one macromolecule in a sense. Examples of a network solid would be something like diamond or silicon dioxide.
A molecular solid is a solid where now we have discrete separate molecules and where those molecules are held together by comparatively weak intermolecular forces. Something like water ice is a good example of this where the individual water molecules are held together by strong covalent bonds (hence why they are called molecules), but the water molecules are held together by intermolecular forces like dipole-dipole and hydrogen bonding which are weaker than a covalent bond. Examples of this would include water already mentioned, but also caffeine and sucrose (table sugar).
If you’re looking how to identify them on a macro scale. Molecular solids generally have lower melting points compared to network solids. Molecular solids are also usually more brittle and soft compared to network solids.
Hope that helps.(8 votes)
- Could somebody please answer this question with a list of the most common covalent network solids? I know diamond and quartz are the most well-known, but what else is there?(2 votes)
- I mean there's quite a lot of them so it depends on how many you actually want to know about. Aside from diamond and quartz the more well-known ones include graphite, graphene, and silicon carbide. Graphite and graphene are allotropes of carbon, or different forms of carbon, alongside diamond.
Graphene is a single sheet of carbon hexagons. Graphite is the material in pencil 'lead' and is actually just lots of layers of graphene stacked on top of each other which can slide off easily as you write. Silicon carbide is a material commonly used as a semiconductor.
Hope that helps.(6 votes)
- Why do some materials sublime?(1 vote)
- Sublimation occurs with chemicals whose triple points are above normal atmospheric pressure.
It's easiest to view these things on a pressure-temperature graph (or PT graph), since physical states depend on both temperature and pressure. A PT graph charts the physicals states a chemical is at as we vary the temperature and pressure and it'll carve out regions in the graph where we would find the chemical as solid, liquid, and gas. Basically it'll show us a set of temperatures and pressure which will give us a certain physical state.
The triple point on a PT graph is the exact pressure and temperature where the solid, liquid, and gas regions intersect and our chemical can exist as all three states simultaneously. It's also the lowest pressure in which a chemical can exist as a liquid.
Normal atmospheric pressure is ~ 101 kPa which on a PT graph would look like a straight horizontal line. If you followed that line you would be able to spot the temperature ranges where a chemical would exist in a particular physical state. For a chemical whose triple point pressure is below normal atmospheric pressure (like water for example) then this horizontal line passes through the liquid region and that chemical would have an intermediate liquid phase and not sublimate. But if the chemical's triple point is above normal atmospheric pressure (like carbon dioxide) then the horizontal line does not pass through a liquid region signifying that we would not have a liquid intermediate and instead sublimate from solid to gas. Or simply the liquid phase requires a higher pressure to exist at than is available at normal atmospheric conditions.
As for why exactly the triple points are at the pressures they're at for certain chemicals, that's beyond my knowledge.
Hope that helps.(6 votes)
- we know that strength of ionic bonds is greater Nand that of covalent bond but still it is observed that the melting point and hardness of covalent solids is greater than the ionic solids. What is the reason behind it??(1 vote)
- Well I wouldn't say that your premise is entirely correct; that ionic bonds are greater (assume you mean stronger) than covalent bonds. There are certainly quite a lot of ionic compounds which can be said to experience stronger bonding than many covalent bonds, but the reverse is true too. We can quantify the strength of a bond by how much energy it takes to breaks the bonds; the more energy required to break the bond the stronger it is. For ionic compounds we can use lattice energy and for covalent compounds we can use bond dissociation enthalpy.
An ionic compound such as sodium iodide has a lattice energy of 704 kJ/mol, while the carbon-oxygen triple bond of carbon monoxide has a bond enthalpy of 1077 kJ/mol. And of course I can list ionic compounds with lattice energies higher than certain covalent bonds, but we can see that one type of bonding is not always stronger than the other.
It's important to note that the type of environment the compound is in also plays a role. For instance many ionic compounds in polar solvents like water dissociate into their ions while many covalent compounds can exist in water without their bonds breaking.
Also I wouldn't use melting point to compare the strength of bonds. Melting a covalent compound doesn't break the covalent bonds of molecular solids, rather it breaks the weak noncovalent bonds like London dispersion forces holding them together. Hardness too doesn't really show bond strength, rather it just exploits any deformities in the solid structure from say a missing atom in the solid.
Hope that helps.(4 votes)
- Is graphite an exception? Because it can conduct electricity and it's a covalent network solid.(1 vote)
- Electrical conductivity in network solids is variable, some are good conductors while others are not so good. It depends on the type of bonding the network solid uses.
In diamond, carbon is bound to other carbons in tetrahedrons using sigma bonds which cause the electrons to stay localized, or do not move around the solid. In graphite, the sheets are made of interlocking hexagon structures with have pi electrons which can be delocalized like the electrons in a metallic solid. So graphite is a good conductor within a sheet, but not between sheets because they only have intermolecular forces holding the sheets together and the pi electrons from one sheet do not cross over into another. This is also why graphene is a good electrical conductor since it’s simply a single sheet of these repeating hexagons.
Hope that helps.(3 votes)
- Silicon is a Metalloid, So is it possible for Metalloids to form covalent bonds?(1 vote)
- Metalloids are interesting in that since they have properties somewhere between metals and nonmetals, they can engage in all sorts of bonding including covalent bonding. The exact type of the bond depends on what silicon is bonded to though.
Hope that helps.(2 votes)
- As in the about, why is graphite 2D and diamond not? I am getting a variety of answers while I am searching for it.(1 vote)
- [Instructor] So we've already talked about multiple types of solids. We've talked about ionic solids. That's formed when you have ions that are attracted to each other, and they form these lattice structures. We have seen metallic solids, and we've thought about them as these positive ions in this sea of negatively charged electrons. And we've also seen molecular solids, which is formed from individual molecules being attracted to each other through intermolecular forces. Now, what's different about covalent network solids is that they're entire networks formed by covalent bonds. What we see here, for example, is a network of silicons and carbons, and this is silicon carbide right over here. And now, some of you might thinking, haven't we already seen covalent bonds involved in a solid before, for example, in molecular solids? And this right over here is an example of a molecular solid that we studied in that video. You have the molecules, which are made up of atoms bonded with covalent bonds. But the reason why they form a solid is because the molecules are attracted to each other through intermolecular forces. And if you wanted to melt this molecular solid, you have to essentially overcome these intermolecular forces. Well, in a covalent network solid, the solid, to a large degree, is made up of these covalent bonds. And if you wanted to melt this somehow, you would have to overcome these covalent bonds, which, generally speaking, are stronger than these intermolecular forces. And so you can imagine, covalent network solids are going to have higher melting points. You also don't see a sea of electrons here. So unlike metallic solids, they're not going to be good conductors of electricity. But just to understand this point a little bit more clearly, let's look at some more covalent network solids. So what you see here on the left, you might recognize as a diamond. A diamond is just a bunch of carbons covalently bonded to each other, and this is the structure of how these carbons are bonded. And as you might already know, diamonds are the hardest solid that we know of. These covalent bonds, the way that they are structured, can take a lot of stress, a lot of pushing and pulling. It's very hard to break it. Now, what's interesting is that same carbon can form different types of covalent network solids. For example, this right over here is graphite, and graphite is probably something you're quite familiar with. When you write with a pencil, you're essentially scraping graphite onto that piece of paper. And so this is what graphite looks like. It's these covalent network sheets, and each of these sheets actually are attracted each other through intermolecular forces. And that's why it's easy to scrape it, because these sheets can slide past each other. But if you really wanted to melt graphite, you would have to break these covalent bonds. And so you can imagine, to overcome the covalent bonds and melt, say, diamond or graphite, it takes a very, very high temperature. Graphite, for example, sublimes at 3,642 degrees Celsius. The silicon carbide that we looked at at the beginning of this video, it decomposes at 2,830 degrees Celsius. This right over here is a picture of quartz, which is a very common form of silicon dioxide, another covalent network solid, and this has a melting point of 1,722 degrees Celsius. So the big takeaway over the last several videos is there's many different ways of forming a solid. It could be with ions, it could be with metals, it could be with molecules that are attracted to each other with intermolecular forces, or you could have a network of atoms formed with covalent bonds.