9
$\begingroup$

I'm not quite sure I understand this. My question is:

Calcium chloride is a salt used widely to melt ice on sidewalks and roads. Explain why one mole of $\ce{CaCl2}$ would be more effective than one mole of $\ce{NaCl}$.

So from everything I've learned, $\ce{NaCl}$ would be more effective for the melting, because calcium chloride dissociates into three ions and sodium chloride dissociates into two, so that would make the boiling point of water with calcium chloride higher. Therefore, it shouldn't be used to melt ice over sodium chloride, as it would make the water have a higher boiling point so wouldn't melt it as well!

$\endgroup$
  • 1
    $\begingroup$ It's a little weird that you're trying to justify changes in the melting point by looking at the effects of dissolved substances on the boiling point. If you have a book around, take a quick look again at freezing point depression (or cryoscopic effect) and compare it with the boiling point elevation (or ebulioscopic effect). In particular, notice how the equations are very similar. What does this similarity imply? $\endgroup$ – Nicolau Saker Neto May 17 '13 at 1:00
12
$\begingroup$

That one mole of $\ce{CaCl2}$ is more effective in melting ice than one mole of $\ce{NaCl}$ is explained by the van 't Hoff factor. The freezing point depression of a solution is calculated by $$\Delta T=K_fbi$$ where $K_f$ is a cryogenic constant which is specific to the solvent, $b$ is the molal concentration of the solute, and $i$ is the van 't Hoff factor, which indicates number of solute particles. In this case, it is the number of ions produced upon dissociation. Because $K_f$ and $b$ are constant, and $i=3$ for $\ce{CaCl2 \rightarrow Ca^2+ + 2 Cl-}$ and $i = 2$ for $\ce{NaCl \rightarrow Na+ + Cl-}$ the freezing point depression is greater for $\ce{CaCl2}$.

The reason $\ce{CaCl2}$ is widely used as melting agent may have more to do with keeping the concrete intact. Concrete is composed of many calcium-containing species. When water flows over the concrete, the calcium can be leeched out, and the concrete becomes brittle. The higher the calcium content in the water, the less leeching occurs, and the concrete remains structurally sound.

$\endgroup$
  • 3
    $\begingroup$ Who cares about moles? One mole of CaCl2 weighs almost as much as two moles of NaCl. So even taking van't Hoff factor into account, CaCl2 should be less effective per kilogram. It is the higher solubility that matters (and makes it more effective, despite the above). $\endgroup$ – Ivan Neretin Jul 7 '16 at 9:15
  • $\begingroup$ Well, the question cited specifically asks for a mol:mol comparison, for one. $\endgroup$ – buckminst Jun 28 '17 at 5:54
4
$\begingroup$

Calcium chloride produces an exothermic reaction when mixed with water or when placed on top of ice, which means that it literally heats up when it comes in contact with ice. Sodium chloride does not produce this exothermic reaction.

$\endgroup$
  • 4
    $\begingroup$ More specifically, you're talking about the enthalpy change of solution. For NaCl, this quantity is slightly positive, about 3.87 kJ/mol, meaning for each mole of NaCl solvated in water, 3.87 kJ of heat is absorbed from the surroundings, making the water colder. CaCl2 has a large, negative heat of solution, seemingly in the range of that of NaOH, thus releasing heat that contributes to melting. $\endgroup$ – Brian Oct 10 '13 at 14:54
  • 1
    $\begingroup$ Exothermic reaction means next to nothing, when you are using it in the street. All that heat vanishes instantly into the environment. $\endgroup$ – Ivan Neretin Jul 7 '16 at 9:18
2
$\begingroup$

In addition to the colligative and thermochemical properties, there are practical considerations as to why calcium chloride or sodium chloride would be used to melt ice in a public area. Calcium chloride will make a surface slippery under reasonably cold ($<0~^\circ\mathrm{F}$) conditions presumably due to the hygroscopic nature of the salt. On the other hand, sodium chloride is fairly corrosive and can damage concrete and vegetation. I'm not a fan of sending visitors to other Q&A sites, but this answer has a nice table of the pros and cons of several salts used in melting ice.

$\endgroup$
  • 1
    $\begingroup$ If you use NaCl on concrete, you will have a water softener where you don't want it. It is one thing to remove calcium from hard water but quite enough to remove calcium from concrete. $\endgroup$ – user55119 Mar 13 '18 at 18:22
-4
$\begingroup$

The answer is simple. One of the unique properties of any salt is that it completely breaks down when dissolved in water. By that I mean each of the atoms clings to water molecules, preventing them from using their natural ionic bonds to each other. For every NaCl molecule, you can occupy 2 water molecules. For every CaCl2 molecule, you can occupy 3 water molecules. This makes it, Mol for Mol more effective an ice melt.

$\endgroup$
  • $\begingroup$ The solvation sphere is much larger than just one molecule of water per ion. $\endgroup$ – Jan Nov 29 '16 at 0:05

protected by orthocresol Nov 28 '16 at 18:23

Thank you for your interest in this question. Because it has attracted low-quality or spam answers that had to be removed, posting an answer now requires 10 reputation on this site (the association bonus does not count).

Would you like to answer one of these unanswered questions instead?

Not the answer you're looking for? Browse other questions tagged or ask your own question.