Mixed Melting Point of Two Pure Solid Crystals

[Mike Dacre]

I understand why an impure solid will have a decreased melting point, however, in my recent organic chemistry lab section I mixed two finely ground pure solids together and observed that the mixture melted at a lower temperature than either pure solid.

I understand this in principle, but whenever mixed melting point theory is taught, it is taught from the perspective that impurities reduce the strength of a crystal structure. However, in my experiment I had two *solid* pure substances, mixed together. The crystal structure could not possibly be disrupted, as the compounds had never mixed in liquid form. So my question is: how does that work? Is it just the case that a small amount of one or both of the pure substances is melting at the edges of the crystals, and that when this happens, this tiny amount of impure liquid can dissolve more of the two pure solids, and that this is what causes the melting point depression?

 

[Bystander]

That's one way to look at it. Strictly? Recall Gibbs' Phase Rule and realize you've got a vapor phase of A and of B in equilibrium with solid A and B, giving you a mixed solid.

 

[Mike Dacre]

That's one way to look at it. Strictly? Recall Gibbs' Phase Rule and realize you've got a vapor phase of A and of B in equilibrium with solid A and B, giving you a mixed solid.

Hi Bystander, thank you for your help but I would like a little extra clarification if possible. Unfortunately, the course I am taking has a separate lab and lecture that aren't always synchronized and I have not learned about Gibb's Phase Rule in class yet. I just went to YouTube and Wikipedia university and I think I understand the phase rule now, but I am not confident that I understand how it relates to this problem.

Here is what I think is happening from your answer: for both A and B, there is a dynamic equilibrium between the solid and vapor phases, meaning that there is deposition of a mixture of A and B on all of the surfaces of the previously pure crystals, which then melt at a lower temperature. As they melt, the mixed liquid is able to dissolve the remaining pure solids, resulting in a pure molten mixture at a temperature well below the melting point of either pure solid. Does that make sense? Or is it that the mixed vapor phase is condensing into a mixed liquid? I am discounting that possibility because I assume that the intermolecular forces that dominate in the condensed phases play only a very minor role in the gas phase, and thus shouldn't play a large role. However, even as I type that I am doubting myself, is that the case?

I assume there are also entropic considerations here, in that the mixed liquid has greater entropy than pure liquids (for ideal mixtures) and thus is generally favored. Is that a safe assumption? Or are so few liquids ideal that assuming anything about entropy is foolish?

 

[Bystander]

As they melt, the mixed liquid is able to dissolve the remaining pure solids, resulting in a pure molten mixture at a temperature well below the melting point of either pure solid.

Pretty much nails it --- the rest of your understanding/discussion looks solid as well.

 

[DrDu]

I think the main mechanism is the contact between the crystals. The crystal surfaces are less stable than the bulk of the crystal and are already liquid at temperatures well below the melting point. The liquid boundaries of touching crystals will mix which reduces the melting point further.

 

[snorkack]

Here is what I think is happening from your answer: for both A and B, there is a dynamic equilibrium between the solid and vapor phases, meaning that there is deposition of a mixture of A and B on all of the surfaces of the previously pure crystals, which then melt at a lower temperature. As they melt, the mixed liquid is able to dissolve the remaining pure solids, resulting in a pure molten mixture at a temperature well below the melting point of either pure solid. Does that make sense? Or is it that the mixed vapor phase is condensing into a mixed liquid? I am discounting that possibility because I assume that the intermolecular forces that dominate in the condensed phases play only a very minor role in the gas phase, and thus shouldn't play a large role.

Presence of significant amounts of vapour is unimportant. Look at a simple case, of bringing into contact solid single crystal sodium chloride with solid single crystal ice.
Saturated solution of sodium chloride in water freezes into eutectic mixture at -21 degrees. Sodium chloride melts at over 800 degrees and boils at over 1400 degrees. At -20 degrees, the vapour pressure of sodium chloride is thus very low.
The mechanism of "surface melting" might count. Still... how quickly is melting initiated at the contact point between ice and salt at -20 degrees?

 

[Mike Dacre]

Great, thank you all for your help.

how quickly is melting initiated at the contact point between ice and salt at -20 degrees

That is actually an interesting kinetics question, and one that I couldn't find an answer to after a little brief searching. If they behave anything like my two unknowns for my organic chemistry lab, they start melting very quickly. Does anyone know the answer? If my proposed mechanism is correct, the rate of melting will be dependent on the surface area of contact between the two crystals, and so finely ground salt added to finely ground ice held at -20 ºC should melt very rapidly (on the order of a few minutes). However, I am not sure what the rate would be if one added coarse salt to the surface of sheet ice, my instinct is that although melting would be initiated quite quickly at the contact point between the surface of the ice and the surface of the salt crystals, so little would melt in that initial contact that the resultant rate would be very slow.

If anyone has anything more quantitative to offer on this, I would be interested to hear.

Thanks again everyone!

 

[256bits]

I mixed two finely ground pure solids together and observed that the mixture melted at a lower temperature than either pure solid.

Not knowing your setup, you might have an experimental error, such as,

There could be interaction between the solids.
There could be presence of water which will lower the melting point. Humidity from the air would be more readily |"absorbed " with a finer powders.
A section of the crucible could hot enough to melt one of the solids, therby enhancing the other solid to dissolve in the liquid.
Your temperature sensing device would most likely record average temperature throughout the sample and not hotter or cooler sections.

I would tend to think that on one trial it would be difficult to conclusively say what is going on, and form a totally conclusive explanation.