24
$\begingroup$

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?

Improve this question
$\endgroup$
6
  • 8
    $\begingroup$ 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, 2015 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, 2015 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, 2015 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, 2015 at 19:37
  • $\begingroup$ Above 2230 C, both are gases. $\endgroup$
    – James Gaidis
    Oct 13, 2020 at 15:55

3 Answers 3

Reset to default
24
$\begingroup$

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.

Improve this answer
$\endgroup$
2
  • $\begingroup$ 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, 2018 at 12:53
  • 5
    $\begingroup$ @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, 2018 at 15:12
7
$\begingroup$

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.

Improve this answer
$\endgroup$
6
  • 1
    $\begingroup$ Silicon dioxide should be nonpolar by symmetry, too, whether in the same structure as CO${}_{2}$ or no. $\endgroup$
    – Jerry Schirmer
    Jul 31, 2015 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, 2015 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, 2015 at 17:47
  • 2
    $\begingroup$ 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, 2015 at 17:51
  • 1
    $\begingroup$ Silicon dioxide is a 3D network not a 2D network. $\endgroup$
    – matt_black
    Jan 10, 2016 at 0:04
3
$\begingroup$

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.

Improve this answer
$\endgroup$

Not the answer you're looking for? Browse other questions tagged or ask your own question.