# How can we tell the particle size of solutes in solutions?

This is a general question in regard to solutes dissolving into solvents. I am interested in the reality of the particle size of the solute during dissolution.

For example, take $\ce{NaCl}$ and dissolve it into water. Are the salt particles, which are now invisible to the eye, truly dissolved? Meaning have they turned into single atom particles of $\ce{Na+}$ and $\ce{Cl-}$ clinging to $\ce{H2O}$ water structures?

Chemical equations seem to imply this. However, when you look at colloids, then particle size matters and can vary. Do ionic salts still have large clusters of their cations/anions upon dissolution?

I am really interested in this topic and would love someone to give me a good explanation.

• Welcome to Chemistry.se! This is not a discussion forum, comments are not meant to stay. Please include all relevant details to your question in the actual question body. Only then will it be useful for future visitors. – Martin - マーチン Jul 19 '17 at 10:13
• Colloidal solutions are not solutions in the sense of the definition, as they are not a homogeneous phase. They can be separated by conventional means, like centrifugation. Apart from this, I think your question is too broad as it seeks information about not only solutions, but also the process of dissolution and is not a good fit for this format. If you limit the scope to only be about NaCl, then that would be more appropriate. – Martin - マーチン Jul 19 '17 at 10:30
• Martin, the question is about ionic compounds. Thanks! – Michael S. Jul 19 '17 at 14:22

I assure you there are many reasons to believe that the ions really are solvated individually. The first thing to look at is the relative strength of interactions between the ions in a lattice and the ions when in a solution. For water, these interactions must be very strong because a number of hydrogen bonds must be broken, the interaction between ion and water must overcome this initial deficit and the loss of some entropy. So, the interaction is surely quite strong.

It is of course possible for two ions which dissolve to get back together and interact with each other. They are oppositely charged ions after all. By looking at the relative size of equilibrium constants for these two processes, however, I think you'll be convinced it is far preferable for the ions to be solvated. For instance, a quick google search will tell you that the $K_{sp}$ of $\ce{NaCl}$ is around $36$. That's an astronomically large solubility product. This means we dump sodium chloride into water until the concentration of both ions is around $6\ M$.

Thus, the easiest experiment to do is to just saturate a solution of water with salt and write the mass of salt added in molarity units. This doesn't really answer your question though because we assumed that the salt existed as two ions and then made a correspondence between that concentration of ions and the concentration of water which each ion is occupying. Plus, we would generally be interested in more dilute solutions than this.

This is why people have used methods which probe the structure of ions in a solution directly. Namely, using x-ray diffraction and neutron scattering experiments, it is possible to determine the local environment around water (or something else solvating an ion) and calculate the hydration number. This is the number of ions, of a specific ion, solvated per molecule of solvent.

Here is a review which discusses x-ray diffraction and neutron scattering experiments and one which discusses coordination numbers of ions solvated in water.

Using these structural methods, which show it is extremely unlikely there are chunks floating around in solution, the following data can be found for hydration numbers of ions in water: $$n_{\ce{Li+}}=4$$ $$n_{\ce{Na+}}=5$$ $$n_{\ce{K+}}=6$$

Here is also an old review which describes many simple but outdated methods for determining the solvation/hydration numbers of ions in a solvent.

Hopefully some of this is helpful and convincing. The experiments are quite definitive, but if you believe the equilibrium constants, then it is very hard to believe these lattices would remain intact in solution.

Using your example of $\ce{NaCl}$, yes, upon dissolution into water the solid is broken down into individual $\ce{Na+}$ and $\ce{Cl-}$ ions in a process called solvation. Each ion is surrounded by a solvation sphere, or solvation complex, of solvent molecules creating a more stable configuration than undissolved $\ce{NaCl}$ in the presence of water (up to the limit of solubility, about $\pu{260}$ g/L at room temperature).

The following image was taken from the Wikipedia article for solvation:

Here you can see that the positively charged sodium ion is surrounded by the electron-rich end of water molecules, creating a solvation spere (of course this is happening in 3 dimensions, thus a sphere). The same thing happens for the negatively charged chloride ions, except that the more electron-poor end of the water molecules are attracted to the ion.

• So, if you took silver metal and dissolved it into nitric acid, based upon this information you would have single atoms of silver surrounded by a solvation sphere. But, when you reduce it again, it turns into colossal metal particles, not a perfect colloid. – Michael S. Jul 19 '17 at 1:16
• I guess I am asking, how do we know that what is at the core of this solvation sphere is truly a single atom and not multiple? As we know colloids that contain many thousands of atoms can appear dissolved in solution. – Michael S. Jul 19 '17 at 1:22
• The definition of a colloid, in terms of size, is a bit fuzzy. Typically they can be considered to be a few nm up to a micron in size. While this seems small compared to bulk material, a colloid still contains thousands to billions of of atoms in size. They can be perceived by light scattering, they can be seen in optical or electrical microscopes. Electrochemical experiments also show that individual ions are at play, and not colloids. – airhuff Jul 19 '17 at 1:25
• Yes, I don't doubt that ions are indeed ionized. I have played a bit with electrolysis myself and it does follow that the positive cations are drawn to the cathode and the negative anions are drawn to the anode. Does this necessarily prove that these are single ions? Or could there be ion spheres with superficial charges that behave as single ions, but really they are large particles? – Michael S. Jul 19 '17 at 1:28
• Another example is ion chromatography (IC). In IC, individual ions interact with a column of material containing charged sites. If the solution being run through the column contained large "superficially" charged colloids rather than individual ions, the interaction would be negligible. I'm not sure that's a great example, but I'm also not sure I understand what you mean by "ion spheres". Do you mean a colloidal particle with a charge to it? Let me think about this a bit and see if I can come up with a better example. – airhuff Jul 19 '17 at 1:39