QP and chemistry

[lemur]

What does knowledge about quantum states, such as electron shape and wave function characteristics, contribute to understanding of how atoms combine and break down in chemical processes? Are their fundamental revelations about atomic structuring that provide revealing insights into how atoms interact that were not apparent prior to quantum theory developments?

[mississippichem]

What does knowledge about quantum states, such as electron shape and wave function characteristics, contribute to understanding of how atoms combine and break down in chemical processes? Are their fundamental revelations about atomic structuring that provide revealing insights into how atoms interact that were not apparent prior to quantum theory developments?

 

Did somebody say my name?

 

Yes, the advent of quantum mechanics has done wonders for the world of chemistry. Before molecular orbital theory, chemistry was stuck with what was called "valence bond theory" which worked quite well for determining the basic geometries of many molecules, but was very lacking with reference to spectroscopy, and other aspects of the observed behavior of electrons in molecules. Many chemical reactions are in fact the result of electrons reaching certain excited states that allow the formation of new orbitals between molecules i.e. new bonds. My area of interest in chemistry is mostly in the department of physical and quantum chemistry.

 

Now days it's possible to predict expected spectral readings or even reaction outcomes sometimes based on the quantum mechanical properties of molecules. It's funny because there is both a journal of molecular physics and a journal of physical chemistry and both really deal with the same subject matter to a large degree. I read them both.

 

I think it is fairly safe to say that quantum theory effects chemistry just as much as physics, especially in these modern times where computational chemistry is on the rise. Math and computer simulations are cheaper than "wet" chemistry once the initial investment is paid off.

 

If you're looking for a good example that is easily accessible, do some reading about electron resonance in aromatic systems [namely benzene]. Classical atomic physics had absolutely no valid explanation for the observed spectroscopy, structure, and reactivity of benzene and other aromatic compounds. Valence bond theory was also not at all satisfactory for the observed behavior of most transition metal chemistry. There is currently a large movement within physical chemistry to make better relativistic corrections to our current quantum mechanical models of molecules involving elements with very large nuclei, something I thought was worth mentioning.

Edited March 6, 2011 by mississippichem

[alpha2cen]

We can calculate molecular structures and their behavior by using quantum physical tools. One of the tools is DFT(density function theory). This approximate formula is very useful for determining organic compounds structure, protein structure and molecular property estimation. Many software are developed for computing this equation. Recently developed good simulation tool is DFT-MD(molecular dynamics), but this software requires high computing resources.

Edited March 6, 2011 by alpha2cen

[lemur]

If you're looking for a good example that is easily accessible, do some reading about electron resonance in aromatic systems [namely benzene]. Classical atomic physics had absolutely no valid explanation for the observed spectroscopy, structure, and reactivity of benzene and other aromatic compounds. Valence bond theory was also not at all satisfactory for the observed behavior of most transition metal chemistry. There is currently a large movement within physical chemistry to make better relativistic corrections to our current quantum mechanical models of molecules involving elements with very large nuclei, something I thought was worth mentioning.

I'm reading the wiki article on benzene and it says that the molecule is thermodynamically stable because the electrons are delocalized. Does this have to do with geometrical structure distributing the bond force in a way that prevents weak-points, or something like that (I know next to nothing about chemistry but it seems logical that if bonding was distributed in a way that favored some bonds and left others weak, that those would be less stable as energy/temperature increased)? Does it also have something to do with the gaseous behavior of the molecule, i.e. that it's "aromatic" because it easily evaporates into air? If so, is that related to the delocalization of electrons? I.e. are the molecules relatively less dense and therefore prone to evaporation?

[mississippichem]

I'm reading the wiki article on benzene and it says that the molecule is thermodynamically stable because the electrons are delocalized. Does this have to do with geometrical structure distributing the bond force in a way that prevents weak-points, or something like that

 

If you work out the orbital configurations for a ring of six carbon atoms (each one having one attached hydrogen), at the end you'll end up with six left over orbitals that are all degenerate and orthogonal to the plane of the hexagonal ring. Because they are degenerate [same energy] they can add together in a linear fashion (remember these orbitals are differential solutions to an equation). Then you end up with even probability distribution of electrons on both faces of the ring. Here is a rendering of the first few [math] \pi [/math]-energy levels of benzene, the orbitals that are depicted horizontally side by side are degenerate:

 

 

Valence bond theory predicts alternating single and double bonds for benzene, but the molecular orbital treatment shows that it ain't so, and really all the bonds are intermediate between single and double bonds. The word "bond" has less and less real meaning as molecules are analyzed quantum mechanically.

 

Does it also have something to do with the gaseous behavior of the molecule, i.e. that it's "aromatic" because it easily evaporates into air? If so, is that related to the delocalization of electrons? I.e. are the molecules relatively less dense and therefore prone to evaporation?

 

The formal definition of aromatic is any collection of atoms that are bonded together in a planar ring containing 4n+2 electrons [where n is an integer] that are not involved in a [math] \sigma [/math]-bond. The word "aromatic" is an artifact of a very old nomenclature system. All aromatic systems have somewhat "delocalized" [math] \pi [/math]-systems; delocalized is sometimes considered a bad word though because technically, all fermions are delocalized to some extent as per Heisenberg. It just so happens though that many aromatic compounds do have a sweet "aromatic" smell and also happen to be evaporated easily (low boiling point, high vapor-pressure at STP)

Edited March 6, 2011 by mississippichem

[lemur]

The formal definition of aromatic is any collection of atoms that are bonded together in a planar ring containing 4n+2 electrons [where n is an integer] that are not involved in a [math] \sigma [/math]-bond. The word "aromatic" is an artifact of a very old nomenclature system. All aromatic systems have somewhat "delocalized" [math] \pi [/math]-systems; delocalized is sometimes considered a bad word though because technically, all fermions are delocalized to some extent as per Heisenberg. It just so happens though that many aromatic compounds do have a sweet "aromatic" smell and also happen to be evaporated easily (low boiling point, high vapor-pressure at STP)

I took "delocalized" to refer to the distribution of electrostatic force and thus mass throughout the molecule; i.e. a relatively non-dense molecule. I wonder about the relationship between phase-change temperatures and molecular structure. Then I wonder about the relationship between molecular structure and quantum structure. I've never really understood relationships between element traits and the characteristics of the molecular compounds they constitute. It just seems arbitrary and random to me. I would like it if I could see some logic in the relationships between sub-atomic forces, molecular structuring, and substance behaviors. I suppose that's a tall order, even at a well-stocked bar.

 

Valence bond theory predicts alternating single and double bonds for benzene, but the molecular orbital treatment shows that it ain't so, and really all the bonds are intermediate between single and double bonds. The word "bond" has less and less real meaning as molecules are analyzed quantum mechanically.

So "single" and "double" bonds don't directly correspond with some interactional behavior of electrons according to their shapes in certain states?

[Horza2002]

Missiippichem has given you the physical chemists views on delocailsation...so I thought I'd give you the organic chemists version. Just remember that they are exactly the same, just explained in a slightly different way.

 

In terms of conjugated systems, delooclisation implies that the electrons can be spread out of several different atoms throught the bonds that connected with. If you like, in a single bond and unconjugated double bond, the electrons stay between the two atoms that are making the bond (ignoring the small contribution from the Hiensburg uncertainity principle). However, when you make conjugated double bonds (double bonds seperated by a single bond), then the pi-orbitlas are able to overlap and so the electrons can now be "spread" out over several bonds. In fact, conjgation is a major tributing factor to acidity of hydrogens in molecules. If the resulting anion is stabilised by delocalisation, the hydrogen is more acidicThe majority of the mass in a molecule comes from the atoms it is made of since electrons mass is soooo much smaller than that of the nuclei.

 

In the case of benzene, once the orginal molecular formula was determined, they said it was a 6-membered ring with alternating single/double bonds. Since double bonds are shorter than single bonds, Kekule structure, this would lead to an odd shaped ring (see file attatched). But as Missiippichem has already said, the bonds are actually halfway between being a single/double bond. The C-C bonds in benzene are 140pm; a normal C-C single bond is 147pm and a double bond is 135pm.

 

Well, a single bond referes to a pair of electrons where the bonding orbtial is between the two atoms....a double bond has two pair of electrons.....and triple bond has six pairs....but once you add in conjugation, then it all starts getting a bit muddy.

 

P.S. Please ignore the awful spelling....I have only just woken up!

delocalisation.pdf

Edited March 6, 2011 by Horza2002

[mississippichem]

I took "delocalized" to refer to the distribution of electrostatic force and thus mass throughout the molecule; i.e. a relatively non-dense molecule. I wonder about the relationship between phase-change temperatures and molecular structure. Then I wonder about the relationship between molecular structure and quantum structure. I've never really understood relationships between element traits and the characteristics of the molecular compounds they constitute. It just seems arbitrary and random to me. I would like it if I could see some logic in the relationships between sub-atomic forces, molecular structuring, and substance behaviors. I suppose that's a tall order, even at a well-stocked bar.

 

I used to wonder the same thing. Next thing you know I have a degree in chemistry

 

Yes that is quite a tall order. I've spent the last four years of my life studying such things, and will probably spend the next six becoming an authority.

 

So "single" and "double" bonds don't directly correspond with some interactional behavior of electrons according to their shapes in certain states?

 

"single bonds" and "double bonds" are really just approximations we use to describe molecular orbitals between nuclei. Each bond type roughly corresponds to a certain symmetry of wave functions. The most concrete thing about bond order is that like Horza2002 said:

 

Well, a single bond referes to a pair of electrons where the bonding orbtial is between the two atoms....a double bond has two pair of electrons.....and triple bond has six pairs....but once you add in conjugation, then it all starts getting a bit muddy.

 

The full descriptions of bonds is quite cumbersome and the nomenclature goes something like this:

 

[math] 2p2p\pi E_{2g} [/math]

 

The above describes what many would call a double bond between two carbon atoms.

[lemur]

In terms of conjugated systems, delooclisation implies that the electrons can be spread out of several different atoms throught the bonds that connected with. If you like, in a single bond and unconjugated double bond, the electrons stay between the two atoms that are making the bond (ignoring the small contribution from the Hiensburg uncertainity principle). However, when you make conjugated double bonds (double bonds seperated by a single bond), then the pi-orbitlas are able to overlap and so the electrons can now be "spread" out over several bonds.

I think I read about this in another book as well. I think I also remember it saying that delocalization of the bond in this way strengthens the overall molecule by making it more intradependent (or something like that, that's my word).

 

 

In fact, conjgation is a major tributing factor to acidity of hydrogens in molecules. If the resulting anion is stabilised by delocalisation, the hydrogen is more acidic.

Because acidity is caused by bonding-behavior tendency due to electron configuration?

 

The majority of the mass in a molecule comes from the atoms it is made of since electrons mass is soooo much smaller than that of the nuclei.

I know that electrons have insignificant mass compared to the nuclei. What I meant was that the structuring of the electrons as they configure molecularly would seem to influence the distribution of nuclear mass within the molecule. Then it seemed like this might determine the way the molecule behaves and its characteristics as a substance.

 

The C-C bonds in benzene are 140pm; a normal C-C single bond is 147pm and a double bond is 135pm.

It sounds like the relationship between bond-strength and length has to do with the tendency of force to increase with proximity, such as in the case of electric charge voltage, magnetic attraction, gravity, etc. This would seem to cause molecules to be denser where stronger bonds are found and less dense where weaker bonds are, unless the bonds were arranged in a way that prevented stronger bonded atoms to be direct neighbors. Bond-strength also seems like it would influence the way that molecules break down when heated, e.g. because stronger bonds would withstand more heating. Is this intuitive reasoning leading me away from factual findings, though?

 

 

 

 

 

[Horza2002]

The acidity of a given proton is dependant on how stable the resulting anion is. For example, carboxylic acids are fairly acidic because the resulting anion is delocalised onto two oxygen atoms (see the attached file above). So essentially, each oxygen atom can be though of has have 1/2 a negative charge. In the case of the green polyene in the file, the deolcalised anion is spread onto six carbons atoms so essentially each as 1/6 a ngetive charge.

 

Another way to look at the bond strenth/length is that as you increase the bond order (single to double to triple....etc) then you are putting more negative charge between the two positive nuclie of the bonds and so they attract each other more and is therefore stronger...you could also say they are stronger because, to break the molecule apart, you have now got two bonds to break. Certianly in organic chemistry, the vast majority or bonds are roughly the same length (there are a few exceptions), so on the whole, one molecules "density" is about equal to that of another...although in chemistry, the electron denisty of a molcules is far more important than where the nuclie are.

 

The bond strength does very much depend on the reactivity of compounds. That is why peroxides are explosive, the oxygen-oxygen single bond is pretty weak so if you heat it up too much, it breaks. Silicon-fluorine bonds (the strongest single bonds there is) are very stable and are basically unreactive. This also translates to higher bond orders as well....nitrogen gas has a triple bond between each of the atoms and is therefore not very reactive to other species (thats not to say it is completely resistant)

[lemur]

The acidity of a given proton is dependant on how stable the resulting anion is. For example, carboxylic acids are fairly acidic because the resulting anion is delocalised onto two oxygen atoms (see the attached file above). So essentially, each oxygen atom can be though of has have 1/2 a negative charge. In the case of the green polyene in the file, the deolcalised anion is spread onto six carbons atoms so essentially each as 1/6 a ngetive charge.

So acidity is due to a concentration of electric charge as a result of ionization the distribution of the "charge surplus" throughout the molecule?

 

Another way to look at the bond strenth/length is that as you increase the bond order (single to double to triple....etc) then you are putting more negative charge between the two positive nuclie of the bonds and so they attract each other more and is therefore stronger...you could also say they are stronger because, to break the molecule apart, you have now got two bonds to break. Certianly in organic chemistry, the vast majority or bonds are roughly the same length (there are a few exceptions), so on the whole, one molecules "density" is about equal to that of another...although in chemistry, the electron denisty of a molcules is far more important than where the nuclie are.

Ok, so what determines the density of a given substance then? Not the size or shape of the atom/molecule? Is one mole of anything always the same volume in the same phase-state, just with different masses due to the atomic weight? Or is their something about the way the particles relate that cause them to form a certain volume?

 

Also, what causes atoms/molecules to be crystaline or ductile? I've read a little about lattice-structuring but I don't understand how this emerges from the (sub)atomic level. Are the substance and molecular levels related in some way?

 

The bond strength does very much depend on the reactivity of compounds. That is why peroxides are explosive, the oxygen-oxygen single bond is pretty weak so if you heat it up too much, it breaks. Silicon-fluorine bonds (the strongest single bonds there is) are very stable and are basically unreactive. This also translates to higher bond orders as well....nitrogen gas has a triple bond between each of the atoms and is therefore not very reactive to other species (thats not to say it is completely resistant)

I've read this about nitrogen and it makes sense to me. I don't know why weak bond strength would make something explosive, though. I would think it would just make it break down at a lower temperature. I would think stronger bonds result in more explosive bond-breakage because more energy could build up before the bond broke to release it. Isn't this why nitrogen is used in explosives? This is somewhat of a tangent, but it is also interesting to me that nitrogen is so useful for plant-growth. I wonder if this is also related to bond-strength.

 

 

[Horza2002]

Acidity depends on how stable the resulting anion is (and other factors aswell). If the density of electrical charge (for lack of a better phrase) is low in the resulting anion, then the proton will be fairly acidic.

 

The density of a substance does depend on the shape of the molecules to and extent; but it is not the only factor. If the compound is very branched, then it'l have a lower density because the molecules can't get that close together. Another contributing factor will be the strength of intermolecular bonds (Van der Waals, ionic interactions, hydrogen bonds). If these forces are strong, then the molecules will tend to get closer together...having said that, the lots of organic compounds have a fairly similar density. While I've been thinking about how to explain this, I have noticed that alot of the reasons are similar to those that affect boiling/melting point actually.

 

Having weak bonds means that it takes very little to break them..and the bonds that are made are even more stable. So you can get essentially chain reactions...a weak bond breaks and releases alot of energy making the new bonds. This energy that is released can then break more of the weak bonds and so on and so on. But also remember that an explosion is a rapid expansion of a gas...as hydrogen peroxide decomposes, it releases water and oxygen gas. The rapid production if the gas is what actually makes the explosion.

 

Nitrogen the element is used in explosives not nitrogen gas. A lot of explosvies combine the two factors I just outlined above. Lots of weak bonds are replaced by extremely strong nitrogen-nitrogen triple bonds which releases a lot of energy that speeds up the cleavage of other weak bonds...the side product is also nitrogen gas that then expands rapidly giving you the explosion. Azide (N₃^-, N^-=N⁺=N^-) used to be used in airbags. On impact, the azide decomposes to release lots of nitrogen gas that fills the air bag.

Edited March 7, 2011 by Horza2002

[lemur]

Another contributing factor will be the strength of intermolecular bonds (Van der Waals, ionic interactions, hydrogen bonds). If these forces are strong, then the molecules will tend to get closer together...having said that, the lots of organic compounds have a fairly similar density. While I've been thinking about how to explain this, I have noticed that alot of the reasons are similar to those that affect boiling/melting point actually.

Would surface-tension in water be an example of intermolecular bonding that sets water's boiling point at a certain level, for example? Or if this is false, could you give an example for reference? Are all liquid/solid densities the product of intermolecular bonding?

 

Nitrogen the element is used in explosives not nitrogen gas. A lot of explosvies combine the two factors I just outlined above. Lots of weak bonds are replaced by extremely strong nitrogen-nitrogen triple bonds which releases a lot of energy that speeds up the cleavage of other weak bonds...the side product is also nitrogen gas that then expands rapidly giving you the explosion. Azide (N₃^-, N^-=N⁺=N^-) used to be used in airbags. On impact, the azide decomposes to release lots of nitrogen gas that fills the air bag.

Wouldn't the triple-bond of nitrogen gas also produce more energy as other molecules fragment and recombine into N2? Also, does the triple bond structure result in stronger electron-exchanges during the chemical reaction, since I'm guessing more bond-energy has to be displaced elsewhere to generate the resultant nitrogen triple bonds? Sorry if I'm speculating too wildly, but it seems like there is logic to this.

[mississippichem]

Would surface-tension in water be an example of intermolecular bonding that sets water's boiling point at a certain level, for example? Or if this is false, could you give an example for reference? Are all liquid/solid densities the product of intermolecular bonding?

 

Water's high boiling point is the result of strong dipole-dipole interactions and in turn hydrogen bonding.

 

Wouldn't the triple-bond of nitrogen gas also produce more energy as other molecules fragment and recombine into N2? Also, does the triple bond structure result in stronger electron-exchanges during the chemical reaction, since I'm guessing more bond-energy has to be displaced elsewhere to generate the resultant nitrogen triple bonds? Sorry if I'm speculating too wildly, but it seems like there is logic to this.

 

 

The above diagram shows the oxidation states and energetics between various common inorganic nitrogen compounds. This is expressed in terms of volts because we are looking at an electrochemical series here, but don't worry too much about that. The important concept is, that we can "reduce" or "oxidize" nitrogen and there are energetic considerations to deal with. Molecular [ce] N_{2} [/ce] is the most stable species on the diagram, and as a result, transforming any of the other nitrogen compounds into [ce] N_{2} [/ce] will release energy. The further away on the diagram, the more energy released. Don't get confused by the [math] \pm [/math] sign convention either. That just has to do with the electrochemical aspect of the diagram. I posted this diagram so you could see the various energetic distances between these species and compare qualitatively.

[Horza2002]

The high boiling point of water is a direct result of hydrogen bonds between the water molecules; an intermolecular force. These are relatively strong so give water its high surface tension and high boiling point. If you look at the boiling points of the hydrogen chalcogen, water is indeed the odd one out because of the hydrogen bonds. The others in this group are not capable of forming hydrogen bonds and so boil at a much lower temperature.

 

Hydrogen oxide H₂O = +100^oC

Hydrogen sulphide H₂S = -60^oC

Hydrogen selenide H₂Se = -41^oC

Hydrogen telluride H₂Te = -2^oC

Hydrogen polonide H₂Po = Not stable

Hydrogen ununhexiumide H₂Uuh = Not been observed

Edited March 8, 2011 by Horza2002

[lemur]

Water's high boiling point is the result of strong dipole-dipole interactions and in turn hydrogen bonding.

Dipole-dipole interactions sound like magnetism. Horza also mentioned hydrogen bonding but I can't figure out how that relates to surface tension.

 

 

 

 

 

So should I read this to say, for example, that NH4+ has more oxidation potential because its number is relatively low, or the opposite? Is HNO3 stronger somehow because it is high on both scales? What does it mean for H2PO4 to have a high oxidation state and low voltage potential (or whatever it is the left axis describes. I appreciate you posting this but I'm afraid I'm having more trouble reading it than you probably thought would be the case.

[Horza2002]

The surface tension is a property of a liquid to minimise the total surface area. This arises because water molecules on the surface have the intermolecular on one side of them pulling them towards the bulk. The hydrogen bond pull the molecules on the sirface towards the bulk

[Horza2002]

The surface tension is a property of a liquid to minimise the total surface area. This arises because water molecules on the surface have the intermolecular on one side of them pulling them towards the bulk. The hydrogen bond pull the molecules on the sirface towards the bulk

[lemur]

The surface tension is a property of a liquid to minimise the total surface area. This arises because water molecules on the surface have the intermolecular on one side of them pulling them towards the bulk. The hydrogen bond pull the molecules on the sirface towards the bulk

Ok, but do you have a simple way of explaining how this works at the level of (sub)atomic forces?

[Horza2002]

Ok, within a water molecule, you have oxygen and hydrogen. Due to the relatively high effective nuclear charge of the oxygen nulcie, electrons around oxygen are more stable as they are closer to the nucleus. This means oxygen has a tendancy to pull electrons in a bond towards itself; its electronegative. Hydrogen, on the other hand, is not electronegative and so the electrons in the H-O bond are shifted towards the oxygen. This results in a partial positive charge on the hydrogens and a negative charge on the oxygen.

 

When there are lots of water molecules around, this difference in charge results in the oxygen of one water molecule being attratced to the hydrogens on another water molecule. Since oxygen has two lone pairs, it can actually form a hydrogen bond with two other water molecules hydrogens. As a result, bulk water molecules (i.e. not on the surface) form tetrahedral structures with the hydrogen bonds holding everything in place. This is a very stable arrangement; the central molecule has four favourable interactions with four other water molecules and so is lower in energy than it would be on its own. Effectively, there is no net force on the molecules, they are being pulled by the same strength force from all directions.

 

However, at the surface of the water, there are only water molecules on one side of the outer molecule. As a result, they are pulled towards to the bulk. Any water molecule on the surface will be less stable than being in the bulk because it can't have four hydrogen bonds around it. This molecule therefore higher in energy than those in the bulk. The system therefore attempts to reduce the overal energy of the system by making liquid have the lowest surface area.

[lemur]

So the relatively high strength of the oxygen in the water molecule results in concentration of positive charge which draws electrons away from the hydrogens somewhat and causes the oxygen side of the molecule to bond intermolecularly with other water molecules through opposite-charge attraction with the hydrogens in those molecules. So the water molecules are electrostatically attracted to each other because their charge is not homogenous throughout the molecule, which makes them behave almost like a magnet (or maybe it would make more sense to compare it to a pair of ions, one negative and one positive). So is this also what causes some kinds of molecules to dissolve in some solvents and not in others? For example, when salt or alcohol dissolves in water, is this due to dipole-dipole attraction between the different molecules? Does oil resist mixing with water because it doesn't have dipole attraction? And the amount of dipole-dipole attraction differs according to the relative weights and structure of the atoms in the molecules along with the distribution of electrons in the specific types of bonds for the molecule, etc.?

 

This is very interesting stuff, btw. I've always thought of chemistry as arbitrary formulas and rules, but this is starting to make it seem like atomic mechanics at the inter-atomic level.

Edited March 9, 2011 by lemur

[mississippichem]

So the relatively high strength of the oxygen in the water molecule results in concentration of positive charge which draws electrons away from the hydrogens somewhat and causes the oxygen side of the molecule to bond intermolecularly with other water molecules through opposite-charge attraction with the hydrogens in those molecules. So the water molecules are electrostatically attracted to each other because their charge is not homogenous throughout the molecule, which makes them behave almost like a magnet (or maybe it would make more sense to compare it to a pair of ions, one negative and one positive). So is this also what causes some kinds of molecules to dissolve in some solvents and not in others? For example, when salt or alcohol dissolves in water, is this due to dipole-dipole attraction between the different molecules? Does oil resist mixing with water because it doesn't have dipole attraction? And the amount of dipole-dipole attraction differs according to the relative weights and structure of the atoms in the molecules along with the distribution of electrons in the specific types of bonds for the molecule, etc.?

 

Well, everything is soluble in everything to some degree, the devil is in the details. Your interpretation is correct enough though. I think you've caught the jest.

Edited March 9, 2011 by mississippichem

[lemur]

Well, everything is soluble in everything to some degree, the devil is in the details. Your interpretation is correct enough though. I think you've caught the jest.

Well, based on the logic of the dipole-dipole attraction through imbalanced distribution of charge between the hydrogens and oxygens in the water molecule, I would guess that water would do a poorer job dissolving molecules with a similar configuration of strong-positive + weak-negative charge, because the positives would repel each other. Are there, on the other hand, molecules with strong negative charge and a weak positive pole that dissolve better in water and worse in each other?

 

 

 

[mississippichem]

Well, based on the logic of the dipole-dipole attraction through imbalanced distribution of charge between the hydrogens and oxygens in the water molecule, I would guess that water would do a poorer job dissolving molecules with a similar configuration of strong-positive + weak-negative charge, because the positives would repel each other. Are there, on the other hand, molecules with strong negative charge and a weak positive pole that dissolve better in water and worse in each other?

 

Actually quite the opposite. Water generally dissolves things that have discrete ions, as well as other polar substances. The water molecules can align itself to solvate positive or negative ions. The sum of the positive charges on the hydrogens is equal in magnitude of the charge on the oxygen atom. Water molecules form a shell around ions in solution accordingly:

 

 

Don't take the picture too literally though, there is really more organization than that but the system is dynamic and constantly shifting, bending and re-coordinating. The water molecules bind the ions strongly, but these interactions are still very weak compared to any covalent bond.

Edited March 9, 2011 by mississippichem

[lemur]

Don't take the picture too literally though, there is really more organization than that but the system is dynamic and constantly shifting, bending and re-coordinating. The water molecules bind the ions strongly, but these interactions are still very weak compared to any covalent bond.

Nice picture, thanks. What is striking to me is that while the intermolecular interactions may be weaker than chemical bonding, they are of a more similar nature than I previously thought. Basically, it sounds like molecules in solutions bond together the way that atoms bond to form a molecule. Apparently, each time bonding occurs at any level, some residual dis-equilibrium is created that results in further bonding potential.

 

The parallels with gravitational potential are really interesting. It's like just as objects fall from positions of higher potential to lower positions within a gravity well, atoms and then molecules fall into each other's electrostatic fields in an entropic process. Does this mean, for example, that some (small amount of) heat is generated when substances dissolve into a solution?