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Molecular vs ionics compounds, H20

by Goodver
Tags: compounds, ionics, molecular
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Goodver
#1
May1-14, 02:35 AM
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I have read previous topics, but still the difference in compounds is not clear to me.

As far as I understood, molecular compounds are made of covalent bonds, while ionic compounds by ionic bonds.

I understand why Mn2O3 is ionic compound, because neutral Mn and O would make only MnO

IN CASE OF H20

So, neutral Oxygen needs to get rid of 2 electrons to complete the shell, Hydrogen needs 1 electron => H20

BUT

What if we deal with H1+ and O2- ions. Then they also can make up the ionic compound then? Because Oxygen now needs 2 positive charges to become neutral and Hydrogen needs 1 charge to become neutral.

So, why H20 is the molecular compound and not ionic?
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Borek
#2
May1-14, 03:32 AM
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Just because we can separate them as ions doesn't mean once they get close to each other, and bond, the bond will be still ionic.
Simon Bridge
#3
May1-14, 03:33 AM
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As far as I understood, molecular compounds are made of covalent bonds, while ionic compounds by ionic bonds.
I've usually thought of this as a bit of hair splitting.

H2O is a molecule because it has covalent bonds. That's it as far as the classification is concerned.

Yes - the ions will attract each other. But, once they get close together, the atoms readily share electrons.

In something like NaCl, the Cl hangs on to it's electron: no sharing, but the attraction is strong enough to hold the atoms together.

Both NaCl and H2O form crystals - but look at the crystal structure.

Goodver
#4
May1-14, 04:44 AM
P: 47
Molecular vs ionics compounds, H20

Ok, is this the reason why molecular compounds tend to be made of non-metals, because non-metals tend to get electrons (high affinity) => more attractive force to form covalent bonds. And ionic compounds made of non-metals because less attractive force which does not cause elements to form covalent bonds, however attractive force made by ions still bond them together?

If so, what determines the "threshold" for elements to bond covalently?

And Al2O3 is ionic compound because Al is a metal => does not tend to attract electrons (in contrast to H20 where H "wanted" to get electrons), Oxygen is non-metal but since, it has completed its shell, does not tend to get more electrons. So, since noone "wants" to get electrons covalent bond does not appear?


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Simon Bridge
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May1-14, 06:34 AM
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If so, what determines the "threshold" for elements to bond covalently?
... whichever has the lower energy.
abitslow
#6
May1-14, 09:11 AM
P: 140
First of all, all electrons of any atom are in orbitals. Second, an ionic bond can be defined as one which is depends only on distance between charges, NOT direction. So, in fact there are very few compounds which are 100% ionic. Most have some directional (covalent) characteristic to the bonds. A great example of this is HCl. In water, it is 100% ionized and most consider it ionic, but in the gas phase it is a covalent molecule. The hydrogen cation is just too small to be purely ionic in condensed phases (fluids, solids), so while I have no direct knowledge of what the HCl crystal structure is (as a solid), I can be confident that it exhibits mostly covalent character.
It is wrong to consider molecules as being un-ionized, just as it is wrong to consider ions as not having covalent bonds. The carbonate anion, CO3(-2) has three covalent bonds and ammonium hydroxide ("ionic") has 5 covalent bonds. On a more advanced level, the charge on an ion or on a molecule can either be localized to one atom or spread over many atoms (via covalent bonds). Carbonate has three equivalent oxygen atoms (in some environments), each having -⅔ charge. This doesn't make sense until you learn a bit about quantum mechanics (simplistically, you can consider these charges to be averages over time, although that is only approximately right). In other environments, the CO3 ion will behave like O=C(O⁻)₂ with two types of oxygen. So, how do you tell? First compare the electronegativity of the atoms, also consider the oxidation state the atoms are in and finally then consider the bonding between atoms as well as their stereochemistry and environment. Except for VERY simple examples, even the experts have to rely on x-ray crystallography (or now days on quantum mechanical calculations which might require supercompeters and a lot of time) to establish the location of charge centers. In other words, even the experts can't just figure it out based on any simple rules or on their experience (although these experts will already KNOW many many of these examples, having learnt them on their way to becoming experts).
Oh, in case you haven't learned about orbitals, only the s (atomic) orbitals are "directionless" (spherically symmetric), so the fact that all electrons are in one orbital or another implies that their "directionality" depends on the orbital they are in. This is heavy duty quantum mechanics, so don't sweat it too much. A beginning student should be learning about the simple cases where these types of facts don't have to be discussed. But only for the most simple metal-nonmetal diatomic compounds is determination of their ionic nature easy. And as I already implied, ionicity depends on pressure, temperature, and other things in the environment, NOT just on the nature of the atoms.
DrDu
#7
May2-14, 01:36 AM
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You can also look at this as follows: Due to the wonderful rules of quantum mechanics, a molecule is never purely ionic or purely covalent but in a superposition of both possibilities. If the ionic contributions dominate, we tend to call the compound ionic, if the covalent contributions dominate, we call it covalent, but there is a continuum of intermediate situations, which are termed polar covalent etc.


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