I was under the impression that chemistry almost exclusively involves valence electrons because there isn't enough energy to strip off electrons located closer to the nucleus.

If that is true, and elements in the same period have similar properties because they have the same number of valence electrons, then why is $\ce{SiO2}$ a solid and $\ce{CO2}$ a gas? Surely a difference in mass by a factor of 2.5 can't be that big a deal.

Is it because of a van der Waals force difference, because silicon has extra electrons which result in compounds being formed from it being more symmetrical than those formed from carbon?

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    $\begingroup$ CO2 is based on double-bonds, O=C=O. So the heart of the question is: "Why can carbon easily form strong double (and even triple) bonds, while silicon is barely capable of forming weak double bonds?" $\endgroup$ – Steve B Jul 31 '15 at 17:16
  • $\begingroup$ Considering that the sublimation point of dry ice is only like $-70{}^{\circ}$ C, a factor of 2.5 of the mass is actually probably plenty to explain this. $\endgroup$ – Jerry Schirmer Jul 31 '15 at 17:32
  • $\begingroup$ Also easy if one is a chemist: Silicium can form $(d-p)\ \pi$ bonds, as it can use its d orbital while carbon can not. Also siliciums radius larger and pi bonds can not be formed efficiently. Long story short: Mainly because of the larger radius of silicium. $\endgroup$ – EpsilonDelta Jul 31 '15 at 17:34
  • $\begingroup$ Do all elements in a column of the periodic table share the same boiling point? No. So why would you expect them to be the same phase of matter at room temperature? $\endgroup$ – pentane Jul 31 '15 at 19:37

It is because of the structure of the $\ce{CO2}$. Two of carbon's valence electrons hybridize into two $sp$ hybrid orbitals. As a result, the molecule is one dimensional with an angle of 180$^\circ$ between bonds and completely non-polar. The $\ce{Si}$, on the other hand, does not form such bonds and the angle is far from 180°, which in conjunction with oxygens high electronegativity makes it quite polar. Thus the interaction between neighboring $\ce{Si}$ and $\ce{O}$ atoms of different $\ce{SiO2}$ molecules is much higher and as a consequence you need much more energy to break the solid, giving it an increased melting point.

The mass (as discussed in various comments) does not play any role here since it is a matter of interaction or forces. The gravitational pull of single atoms or molecules is ridiculously small and never finds any considerations in such calculations (as it should!).

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    $\begingroup$ There are no different molecules in solid SiO2 - it's 3D network! $\endgroup$ – Mithoron Aug 12 '15 at 14:41
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    $\begingroup$ Then interpret the word different in the right way.. $\endgroup$ – t0xic Aug 13 '15 at 17:10
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    $\begingroup$ The point is that the particular molecules of CO2 that form solid CO2 will be released when the CO2 sublimes. However as SiO2(g) condenses into solid SiO2 and you break apart the solid SiO2, you won't get the "same" SiO2(g) molecules. The Si and O lose their molecular associativity in the solid. $\endgroup$ – MaxW Mar 2 '17 at 18:13
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    $\begingroup$ There are no molecules of SiO2. Silica is an infinite network/covalent solid. $\endgroup$ – matt_black Aug 1 '17 at 23:39
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    $\begingroup$ I'm somewhat worried that blatantly incorrect and confusing answers are being accepted as the correct answer. $\endgroup$ – matt_black Sep 6 '17 at 17:28

The reason why carbon dioxide is a gas and silicon dioxide is a solid is because their chemical structures are different.

Carbon dioxide is a linear structure with two double bonds between carbon and oxygen. It is a small molecule and non-polar with only weak bonds between the molecules. Hence it is a gas.

Silicon dioxide is not formed of small molecules. It consists of an infinite array of silicons where each silicon is bonded to four separate oxygens (and each oxygen is shared between two silicons). This creates a strong refractory solid (glass and sand are mostly silicon dioxide aka silica). So the same apparent overall formula doesn't describe the actual structure of the compounds at all. But the structures explain the difference in behaviour.

Of course this doesn't explain why silicon prefers to bond with four oxygens when carbon prefers just two. This is not completely simple and results from the relative bond strengths of carbon-oxygen bonds, carbon-oxygen double bonds and the equivalent bonds for silicon and oxygen. The simple version is that silicon oxygen bonds are strong relative to their double-bond equivalents whereas carbon-oxygen double bonds are strong relative to their single bond equivalents. Or, more precisely, if we could make a carbon-oxygen network solid with the equivalent structure to silica, it would tend to fall apart into carbon dioxide. If we could make silicon dioxide molecules, they would react with the release of energy to give silica.

Deeper explanations would need to look at why the relative strengths of double and single bonds turn out that way, but that would get into molecular quantum mechanics and would not be much more useful as an explanation.

The simplest explanation is the fact that the structures are different.

  • $\begingroup$ Hi, thanks for your answer! In my book (amazon), it is given that "Si cannot form stable bonds with O, hence, Si has to form a 3d lattice. This is the reason why $\ce{CO2}$ is a gas while $\ce{SiO2}$ is a solid" What would you comment about it? Is it correct or not? Thanks! $\endgroup$ – Gaurang Tandon Feb 4 '18 at 12:53
  • $\begingroup$ @GaurangTandon If your quote is exact, then it is wrong. Si is happy forming stable bonds to oxygen: the issue is that collectively single Si-O bonds are far more stable than double bonds. You can't make a small uncharged molecule with just single Si-O bonds so you end up with a network. The issue is the relative stability of single and double bonds. $\endgroup$ – matt_black Feb 4 '18 at 15:12

If you look at the crystal structure of silicon dioxide ($\ce{SiO2}$), you will see it is comprised of tetrahedra (one silicon surrounded by 4 oxygen atoms) and these tetrahedra are in turn connected to each other, and subsequently form a 2-dimensional network which is large enough to make this substance a solid.

Also, the factor of 2.5 in mass makes a big difference, but it is not relevant to this discussion. I would recommend reading any inorganic chemistry book where such topics are discussed in depth. Van der Vaals forces only play significant roles in long-chain carbohydrates, such as fats, and are not really observed in "inorganic" molecules.

$\ce{CO_2}$ molecules cannot form a crystal in the way $\ce{SiO2}$ units do, and I do not think I have to explain why. Also, it has no dipole moment and no van der Waals forces between the molecules.

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    $\begingroup$ Silicon dioxide should be nonpolar by symmetry, too, whether in the same structure as CO${}_{2}$ or no. $\endgroup$ – Jerry Schirmer Jul 31 '15 at 17:34
  • $\begingroup$ It is, but it does not really matter for this question, because as I stated, SiO2 forms large 2-D networks. $\endgroup$ – EpsilonDelta Jul 31 '15 at 17:45
  • $\begingroup$ Sure, but you bring up the fact that CO${}_{2}$ has no dipole moment, which doesn't distinguish it from SiO${}_{2}$. What structure does dry ice take? I would assume it would be similar to that of solid silicon dioxide. I really think the mass is the main difference here. $\endgroup$ – Jerry Schirmer Jul 31 '15 at 17:47
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    $\begingroup$ You are very wrong, it is not even a bit similar to the structure of solid CO2. The mass difference is surely not the main difference, every inorganic textbook will show you. $\endgroup$ – EpsilonDelta Jul 31 '15 at 17:51
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    $\begingroup$ Silicon dioxide is a 3D network not a 2D network. $\endgroup$ – matt_black Jan 10 '16 at 0:04

Elements of the answer are also contained in other answers here but this needs some more teasing out. It has to do with the bond energies of C-O, Si-O, C=O, and Si=O. I don't have the numbers at hand, but work it out and it will point you to carbon being the most stable in the form of $\ce{CO2}$ and silicon the most stable as a tetrahedral network (similar to diamond.)

The reason why $\ce{CO2}$ is a gas can easily be explained using intermolecular bonding principles, which you can look up easily.


Silica ($\ce{SiO_2}$) has a three dimensional structure. It has very strong Si-O bonds and it has a high melting point.

However, in $\ce{CO2}$ (which has a linear shape) there are weak C-O bonds with no dipole moment and it has sp hybridisation. Due to this type of bonding and large gaps in $\ce{CO2}$ molecule, $\ce{CO2}$ is a gas.


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