Silicon dioxide


by Bladibla
Tags: dioxide, silicon
Bladibla
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#1
May5-05, 04:09 PM
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Hey all

Why does Silicon Dioxide have a Giant molecular structure, while Carbon dioxide is made of simple molecules?

They both have 4 electrons on their valence shells, so why wouldn't Silicon Dioxide be a simple linear molecule? (gaseous).
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The Bob
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#2
May5-05, 04:47 PM
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Quote Quote by Bladibla
Hey all

Why does Silicon Dioxide have a Giant molecular structure, while Carbon dioxide is made of simple molecules?

They both have 4 electrons on their valence shells, so why wouldn't Silicon Dioxide be a simple linear molecule? (gaseous).
It is mainly due to the size of the silicon atom in comparison with the carbon atom. The silicon atom is bigger and so there is less room for the oxygen to be attracted for form a double bond so it has to form a lattice ot be stable.

Correct if I am wrong (anyone).

The Bob (2004 )
Gokul43201
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May5-05, 04:50 PM
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This is a tough one, and there may not be any simple intuitive answer. I know that CO2 has been polymerized at high pressures.

I suspect that part of the reason may be the differences in covalent radii. The carbon atom is roughly the same size as an oxygen atom, but a silicon atom is about 50% bigger. This makes it easier to put more O atoms around the Si atom. In the case of CO2 there would likely be large steric energies from trying to force monomers together.

EDIT : Whoops !! Looks like my guess is similar to what The Bob is saying.
The Bob : have you read this somewhere ? If you have a reference or source, that would be useful. My post is merely an educated guess.

DaveC426913
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May5-05, 04:54 PM
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Silicon dioxide


Also, the outer shell of electrons is held much more loosely.


I don't know what you mean by "giant molecular structure", or "simple linear molecule". Do you mean why does it crystallize?
Bladibla
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May5-05, 05:55 PM
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Quote Quote by DaveC426913
Also, the outer shell of electrons is held much more loosely.


I don't know what you mean by "giant molecular structure", or "simple linear molecule". Do you mean why does it crystallize?
Sorry, should have made myself more clear on that. I was trying to ask why Sillicon dioxide looks like this
, while Carbon Dioxide is a gas (RTM).

So yes, you are correct. why does Sillicon dioxide crystallise while CO2 is a gas?
Artermis
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May5-05, 07:42 PM
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Bladibla, am I just paranoid or were you doing the 2004 Chem 13 News Exam? =)

Hehe. I actually didn't know the answer to that question either... CO2 vs. SiO2... I eliminated my way to the answer though; I'm still kind of not understanding of why CO2 is a molecular compound while SiO2 is a covalent network solid... it does spark my knowledge of diamond being the same...
Gokul43201
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May5-05, 08:43 PM
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Quote Quote by Bladibla
Sorry, should have made myself more clear on that. I was trying to ask why Sillicon dioxide looks like this
, while Carbon Dioxide is a gas (RTM).

So yes, you are correct. why does Sillicon dioxide crystallise while CO2 is a gas?
This is a slightly different question from what you had in the OP. You are now talking about the physical state, rather than polymerization or the lack of it.

Your question now reads : Why is SiO2 a solid, while CO2 is a gas at STP ?

In any case, the relative atomic radii are probably still important. But the more dominant reason is likely the electronegativities. Carbon is much more electronegative than Silicon. As a result, CO2 is much less polar than SiO2. So, the interaction between CO2 molecules is much smaller than that between SiO2 units. Hence, CO2 is a gas at temperatures/pressures where SiO2 is a solid.
Bladibla
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May6-05, 11:20 AM
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Quote Quote by Gokul43201
This is a slightly different question from what you had in the OP. You are now talking about the physical state, rather than polymerization or the lack of it.

Your question now reads : Why is SiO2 a solid, while CO2 is a gas at STP ?

In any case, the relative atomic radii are probably still important. But the more dominant reason is likely the electronegativities. Carbon is much more electronegative than Silicon. As a result, CO2 is much less polar than SiO2. So, the interaction between CO2 molecules is much smaller than that between SiO2 units. Hence, CO2 is a gas at temperatures/pressures where SiO2 is a solid.
Damn.. sorry about that. Give me some time to reform that question..
The Bob
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#9
May6-05, 12:35 PM
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Quote Quote by Gokul43201
EDIT : Whoops !! Looks like my guess is similar to what The Bob is saying.
The Bob : have you read this somewhere ? If you have a reference or source, that would be useful. My post is merely an educated guess.
Salters Advacned Chemistry Textbook - Chemical Ideas (Second Edition) - Section 5.2 Pg. 91.

CO2 and SiO2 are both in Group 4 and similar properties of the two would be expected. However, CO2 is a gas (R.T.P.) and SiO2 is a solid (R.T.P.). The physical properties are different due to the bonding. Both are covalent compounds but the size of the carbon atom makes it possible to form double bonds with oxygen. Therefore, CO2 is made of individual molecules. The attraction between the forces are weak (weak intermolecule forces of attraction). Little energy is need to seperate them so it is a gas (R.T.P.).

Silicon atoms are larger than carbon atoms and they normally bond to four oxygens. Quartz (SiO2) is an extended network of SiO4 units, where the central silicon is covalently bonded to four oxygen atoms. Each silicon atom has a half-share of oxygens. Due to this large strucure, it is insoluable in water and has a high melting and boiling points (making it a solid at room temperature) because of the strong covalent bonding throughout Quartz.

Carbon dioxide is an example of a covalent molecular structure.
Silicon (IV) oxide is an example of a covalent network structure.

The Bob (2004 )
GCT
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May6-05, 09:01 PM
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What we seem to be dealing with here are types of intermolecular bonds, all of this should be explained adequately in your text. What they meant by sillicon being "larger" is in reference to the quantum level of its valence electrons; note that pi bonding is due to overlap of individual pi atomic orbitals, it is better achieved with electrons in the same quantum number.

Due to the linear geometry, carbon dioxide has net dipole moment of zero, there won't be much intermolecular attraction, and also I would imagine that sillicon/oxygen interactions are distinctly different in nature than that of carbon and oxygen...I'll have to do a bit of research to figure out the more vital specifics.
Gokul43201
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May7-05, 01:07 AM
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Quote Quote by GCT
What they meant by sillicon being "larger" is in reference to the quantum level of its valence electrons; note that pi bonding is due to overlap of individual pi atomic orbitals, it is better achieved with electrons in the same quantum number.
I have no idea what you are talking about here...
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May7-05, 04:09 PM
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Why does the atomic radii increase as one moves along a column in the periodic table e.g. Li to Na? The valence electrons are in successively higher quantum levels (number), the valence electrons are located farther away from the nucleus. Elements in the same period will be able to pi bond more effectively. It's just something I remember reading about.
The Bob
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May7-05, 05:43 PM
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Also you need to consider going down the group in the case of silicon and carbon. There are more shells between silicon and it's valence shell and carbon's.

The Bob (2004 )


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