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Name: Justice
Status: student
Grade: 6-8
Location: WI
Country: N/A
Date: April 2007

Question:
When you have a mixture of water and salt and the water evaporates, why does the salt crystallize?



Replies:
Hello Justice,

A particular volume of fluid is only capable of having so much of a salt dissolved in solution. So if you dissolve enough salt in a solution there will be a point at which the solution is saturated (cannot have any more salt dissolved) and any further salt poured into it will not dissolve. Let us suppose you have a solution that is saturated, but you do not pour any more salt into it and set it aside to evaporate. As the water evaporates from the solution, there will be too much salt left over in the remaining solution and it will come out (if possible) to form small crystals. As the water continues to evaporate, the salt will deposit on the crystals causing them to grow in size. This is one of the primary ways in which crystals are made.

But why does this produce crystals? Why not just lumps of the salt?

The lowest energy state often corresponds to a regular, closely packed arrangement of atoms or molecules. Just as objects will roll down hill if possible, atoms will attempt to bond in a way that "rolls down hill" which usually corresponds to an ordered arrangement. Suppose you have a small crystal already present. An atom or molecule attempting to come out of solution can try to bind in many different ways and at many different sites. If the molecule attempts to attach in a haphazard way, it is likely to lose contact with the solid. If it happens to find a good location and configuration, it is likely to stay there permanently and become part of the growing solid.

But how do you get the "tiny" solid crystal in the first place for the rest of the salt to deposit onto?

It can happen one of two ways (potentially). The easiest way is if there is already something present in the solution which the salt happens to favor bonding to. This can often be dust or an impurity, but can also (very often) be just a small (too small to see) crystal of the original salt that didn't dissolve. Just because you cannot see any salt crystals, does not mean that they are not already present! These impurities or "seed" crystals act as nucleation sites allowing for a crystal to grow.

However, it can happen even without anything extra. If you have got enough molecules attempting to come out of solution, and enough patience, the molecules will eventually form crystals. It works the same way except it requires a few more steps. You need nearby molecules attempting to come out of solution at the same time to find each other and bond. Most of the time even when this happens the molecules will just lose contact and go back into solution. However, eventually you will get other molecules that will attach to the first pair before they are able to split. And still most of the time, these few molecules will get back into solution. But rarely (and how rarely depends on the molecules, solution, thermodynamics, and more) even more molecules will find the little bundle before it separates. There is a critical size beyond which the little mass crystal becomes stable and will no longer dissolve. You just need to wait long enough for the right conditions and for enough molecules to find each other. These two techniques can then be used to grow the desired kinds of crystals. If you want lots of small crystals, then the "assisted" nucleation is better. If you want fewer, but larger crystals, then the latter approach is better.

best wishes,

Michael S. Pierce
Materials Science Division
Argonne National Laboratory


Justice,

Salt crystallizes because the crystalline form is energetically favored. While it is true that when salt is dissolved in water, it will be ionized, this is only so because the ionized form of salt is stabilized by the interaction with water.

Imagine sodium chloride and the energy it takes to form that. Sodium would require energy in order to change from sodium metal to sodium ions (this energy is often referred to as ionization energy). While there is a release of energy in ionizing chlorine, there is also a required energy in converting chlorine gas (Cl2, the stable form of chlorine) to chlorine atoms. Thus, if the salt does not crystallize, there is an overall increase in the energy of the system - and the rule of thumb is that if the energy of the system increases, then it is more unstable. In the process of crystallization, there is a release of energy (referred to as lattice energy). Thus, when sodium and chlorine gas react to form sodium chloride, there is an over-all release of energy because the gains in energy in ionization (of sodium) and atomization (of chlorine) is outweighed by the loss in energy in ionization (of chlorine) and crystallization (of sodium chloride). If no crystallization occurs, as in the solubilization of the salt, than this reduction in energy is supplied by the interaction with the solvent. If there is no solvent, as in the case when the water is evaporated, than crystallization must take place in order for the system to have a lower energy and become stable.

Greg (Roberto Gregorius)


Salts such as sodium chloride (NaCl, also known as table salt) tend to have very high melting points (around 800C for NaCl) since salt crystals are held together by the strong attraction of atoms with opposite charges. In this case sodium has a full positive charge and the chloride ion has a full negative charge.

In a polar solvent such as water, each solvent molecule has a region which is partially positively charged and a region is partially negatively charged. Salt is dissolved when atoms are released from the crystal and surrounded by solvent molecules, which can orient themselves so that each ion is surrounded by areas of solvent molecules with opposite partial charge.

As water evaporates from a salt solution, salt ions are no longer surrounded by solvent molecules. Therefore the ions will interact with each other, and due to the strong forces resulting from their opposite charge will reform a crystal lattice.

Ethan Greenblatt
Stanford Department of Chemistry



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