Mechanism of successively melting ice with table salt [closed]








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Hi I was reading about how regular sodium chloride in itself can melt ice outside, for example ice on roads during low temperatures. However, I am unsure about the exact details, as I have found some conflicting information, which I would like to clarify here.


When we neglect any temperature change and assume the air as well as the ice are at the same temperature, maybe -2 degrees celsius, what would happen is the following:


When pouring the sodium chloride crystals on the ice surface, because of the room temperature crystals, heat is transferred to the ice, inducing a phase change to water, so that the NaCl dissolves and forms a concentrated solution, additionally the emerging solution should also cool down due to the positive solvation enthalpy of NaCl.


(On this website, it is stated that ice always has a water layer on top of it; to my knowledge, that equilibrium H₂O(s)⇌H₂O(l) only exists at 0 degrees celsius so should not apply here. <https://www.scientificamerican.com/article/why-do-we-put-salt-on-icy/>)


This concentrated solution has experiences freezing point depression (a 1M solution would freeze at around -4 degrees celsius), so that any ice surface in contact with the solution should dissolve, as the ice in equilibrium with the undercooled (negative degrees celsius) solution should not be able to mantain its solid form and also experience a phase transition to water.


(Would this actually also mean, that solid ice would also dissolve by itself in a concentrated solution of sodium chloride at negative degrees celsius?)


This process continues until the sodium chloride solution is too diluted to dissolve any more ice in it.


Would this be the correct explanation of the individual steps of this process? Would other salts execute this process more quickly? (For example salts with positive solvation enthalpy)


Any help or input would really be appreciated, as there is some conflicting information out there.


Cheers!




It is well-known that “Solids do not react together at room temperature over normal time scales and it is necessary to heat them to much higher temperatures, often to 1000 to 1500 °C, in order for the reaction to occur at an appreciable rate.” Ref 1. One reason this is true is that even when molecules of one solid diffuse into another solid, the resultant is still a solid; the only thing observed is a diffusion, and this is so slow that it seldom qualifies as a “reaction”.


But “Thus, we study several reactions that have been recently reported to proceed “in the solid phase” and clearly show that, in most cases, grinding the two solid reactants together results in the formation of a liquid phase.” Ref 2.


To be very accurate, the salt doesn’t melt the ice; a salt *solution* causes ice to melt. If you don’t have a liquid phase to begin with, you won’t get any ice melted.


Consider ice and water in dynamic thermal equilibrium at 0ºC. Water molecules pass between the solid ice and liquid water at equal rates in opposite directions. The ice crystals may rearrange slowly, but the amount of ice and water remain constant. This balance is maintained at 0ºC, but can be upset, for instance, by adding salt.


Combining salt and ice (at the same temperature, below 0ºC), cause two reactions to occur simultaneously: #1. A solution is formed from traces of water, or from grinding. The temperature is reduced slightly because energy is required to dissolve the salt: about 1 kcal/mole of salt. This does not involve melting of the ice, just dissolution of salt by a (tiny bit of) liquid H2O. #2. Water molecules in the ice pass into the salt solution, but fewer pass back to the ice. This means that the ice is melting. The heat required to melt the ice (68.3 kcal/mole of H2O) comes from the salt solution and reduces its temperature much more.


These two reactions result in ice being drawn into the salt solution, increasing its volume at the expense of ice volume, until the ice and solution temperature are reduced and the water concentration is increased to a point where the water molecule exchange rates become equal again.


This process is used in two ways: #1: In the lab, for cooling baths, with excess salt added so that the mixture reaches its lowest temperature (-21ºC Ref 3.) #2: On roads, salt is spread onto a large amount of ice/snow, liquifying it by creating a combination of salt and water which has a melting/freezing point below the ambient temperature. Salt is spread as a fine granule or as a solution for faster action. Other salts, e.g., CaCl2, usually hydrated, and therefore somewhat wet, are preferred for melting ice at lower temperatures.


Ref 1. <https://en.wikipedia.org/wiki/Solid-state_reaction_route>
Ref 2. Gadi Rothenberg, Andrew P. Downie, Colin L. Raston, and Janet L. Scott, J. Am. Chem. Soc., 2001, 123, 36, 8701–8708, <https://doi.org/10.1021/ja0034388>
Ref 3. <https://sciencenotes.org/why-salt-makes-ice-colder-how-cold-ice-gets/>


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