The Medium.com article Mars Phoenix Lander, 10 Years Later shows several remarkable images and discoveries on Mars by the Mars Phoenix Lander circa 2008.

One image (shown below) shows what looks like droplets of liquid water, condensed on the surface of one of the lander's legs.

The article says (emphasis mine):

Shortly after landing, the camera on Phoenix’s robotic arm captured views of blobs of material on one of the landing struts. Over time, these blobs moved, darkened, and coalesced, behaving like droplets of liquid water. The hypothesis here was that these blobs “splashed up” on the struts when the descent thrusters melted the ice exposed upon landing mentioned above.

But if liquid water isn’t stable on the martian surface, how did Phoenix observe liquid water on Mars? The key here lies in salt. If you live anywhere that gets snow, you’re probably familiar with salt as a de-icer for roads, sidewalks, etc. Salt lowers the freezing point of water, allowing it to remain liquid at temperatures lower than that of non-salty water. For example, pure water freezes at 0 °C/32 °F, but ocean saltwater freezes around −2 °C/28.4 °F. While the de-icing salts you get at the hardware store lower the freezing point by a few degrees, more exotic salts can lower the freezing point as much as −70 °C/−89 °F! Phoenix discovered some of these exotic salts in the soil around the lander—in particular, magnesium perchlorate. (note, minor editorial changes have been made)

Question: Which "exotic salt" can lower water's freezing point by 70 °C?

Is it in fact magnesium perchlorate (which was found on Mars) or is it a different salt?

Blobs of possible brine (really salty water) imaged on one of Phoenix’s landing struts shortly after arriving on Mars.

Blobs of possible brine (really salty water) imaged on one of Phoenix’s landing struts shortly after arriving on Mars. Credit: NASA/JPL-Caltech/University of Arizona/Max Planck Institute

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    $\begingroup$ You know why freezing point of water is decreased by using salt? It is because of relative lowering in vapour pressure which depends on number of particles present in solution, it doesn't matter whatever sizes are or whatever the salt is, the necessary conditions are that the substance you are mixing must be non volatile and it must be a solution. So basically the more salt you mix the lower the freezing point is but there is a limit on how much you can mix salt in water. And freezing point is decreased by a few degrees only. -70° is like a dream. the text you are reading might be wrong. $\endgroup$
    – knoftrix
    Commented May 5, 2019 at 2:26
  • $\begingroup$ Also -70 C seems awful cold for a super cooled liquid. $\endgroup$
    – MaxW
    Commented May 5, 2019 at 2:42
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    $\begingroup$ @SauravSingh, vapor pressure has nothing to do with freezing point, though it does affect the boiling point. $\endgroup$ Commented May 5, 2019 at 3:19
  • $\begingroup$ @DrMoishePippik as an aside, vapor pressure may have a role in the appearance of the droplets nonetheless. Assuming this is indeed water, it may have been the heat of the spacecraft that drove water vapor out of the Martian soil, allowing some of it to "deliquesce" on the salt-encrusted leg of the spacecraft. $\endgroup$
    – uhoh
    Commented May 5, 2019 at 3:25
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    $\begingroup$ If "lower water's freezing point by 70 °C" makes its freezing point -70 °C, then "lower water's freezing point by −70 °C", lowering it by a negative amount should raise its freezing point to 70 °C. I think you should remove that "-" from the title and question itself. $\endgroup$ Commented May 6, 2019 at 22:32

3 Answers 3


I recently got a chance to attend a talk by someone who was working on developing analytical instrumentation on Mars. The interesting story is that the initial results by ion-selective electrode was that Mars soil is full of nitrates. Nobody knew on Earth that the nitrate ion selective electrode is far more responsive to perchlorate than nitrate. After learning this, it was an eye opener for analytical chemists! Now they wish to use chromatography rather than electrochemistry. So this was a good lesson for us on Earth.

The perchlorate ion was discovered in 2008 by the nitrate selective electrode. No specific electrode was attached to detect perchlorate, it was rather an accidental discovery. The Science Report makes a footnote "Detection of Perchlorate and the Soluble Chemistry of Martian Soil at the Phoenix Lander Site" (paper: Science 2009, 325 (5936), 64–67)

A Hofmeister anion ISE was intended to monitor nitrate from a $\ce{LiNO3}$ reference electrolyte that was part of the leaching solution, but was ultimately used for perchlorate detection

[Footnote] The relative sensitivity of the Hofmeister series ISE to perchlorate over nitrate is 1000:1, and substantial quantities of perchlorate will overwhelm any other signal. If, as was observed, >1 mM perchlorate accounts for the observed signal, it would require >1000 mM nitrate to produce the same response. This would correspond to more than the mass of the entire sample.

Now that they know it is a perchlorate ion, people did some studies on supercooled brines. See this paper: Toner, J.; Catling, D.; Light, B. The formation of supercooled brines, viscous liquids, and low-temperature perchlorate glasses in aqueous solutions relevant to Mars. Icarus 2014, 233, 36–47 (also available here). They clearly show that if calcium or magnesium perchlorates are slowly cooled, one can get supercooled brines up to -120 Celcius. This is a rather amazing finding. They call it a glassy state.


Your Question: Which "exotic salt" can lower water's freezing point by $\pu{-70 ^\circ C}$?

Here is your "exotic compound" although it is not a salt by definition. It is a base: Aqua ammonia, also called ammoniacal liquor, ammonia liquor, or ammonia water, is produced by dissolving ammonia gas ($\ce{NH3}$) in water. The proper chemical name of aqua ammonia is ammonium hydroxide ($\ce{NH4OH}$), which is in the following equilibrium with water: $$\ce{NH3 + H2O <=> NH4+ + OH-}$$

Ammonia is very soluble in water: According to Wikipedia, its solubility in water is $47\% (w/w)$ at $\pu{0 ^\circ C}$, $31\% (w/w)$ at $\pu{25 ^\circ C}$, and $18\% (w/w)$ at $\pu{50 ^\circ C}$. Therefore it is ideal to cause large freezing point depression since its solubility increases with decreasing temperature. Now, let's see how are the freezing points of aqua ammonia solutions behave with increasing concentrations. The large scale manufacturer of aqua ammonia, Tanner Industries, listed following values of boiling and freezing points of various solutions in its Customer Manual: $$ \begin{array}{ccc} \\\hline \% \ce{NH3} \text{ (by weight)} & \text{Approx. Boiling point} & \text{ Approx. Freezing point} \\\hline 23.52 & \pu{103 ^\circ F}\: (\pu{39.4 ^\circ C}) & \pu{-56 ^\circ F}\: ( \pu{-48.9 ^\circ C})\\ 25.48 & \pu{95 ^\circ F}\: (\pu{35.0 ^\circ C}) & \pu{-69 ^\circ F} \: ( \pu{-56.1 ^\circ C})\\ 27.44 & \pu{88 ^\circ F} \: (\pu{31.1 ^\circ C}) & \pu{-89 ^\circ F}\: ( \pu{-67.2 ^\circ C})\\ 29.40 & \pu{85 ^\circ F}\: (\pu{29.4 ^\circ C}) & \pu{-110 ^\circ F} \: ( \pu{-78.9 ^\circ C})\\ 31.36 & \pu{73 ^\circ F}\: (\pu{22.8 ^\circ C}) & \pu{-123 ^\circ F} \: ( \pu{-86.1 ^\circ C})\\ 33.32 & \pu{66 ^\circ F} \: (\pu{18.9 ^\circ C}) & \pu{-148 ^\circ F}\: ( \pu{-100 ^\circ C})\\\hline \end{array} $$

Accordingly, anything between $29-33\%$ of aqua ammonia solution would do the job.

On the other hand, if you are looking for only a "exotic magic salt," then, don't look too far: Ammonium fluoride ($\ce{NH4F}$) would do the job. The solubility of $\ce{NH4F}$ in $\pu{100 mL}$ of water is listed as $\pu{100 g}$ at $\pu{0 ^\circ C}$ (Wikipedia). That's give you a $\pu{27.0 m}$ solution at $\pu{0 ^\circ C}$. Theoretically, that would freeze at $\pu{-100.4 ^\circ C}$! ($\Delta T = \pu{27.0 m} \times 2 \times \pu{-1.86 ^\circ Cm^{-1}} = \pu{-100.4 ^\circ C}$)

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    $\begingroup$ @uhoh I'd ask the definition of a salt, if $\ce{NH3}$ or any of the $\ce{NH_4X}$'s don't fit the bill... $\endgroup$
    – Stian
    Commented May 5, 2019 at 8:16
  • $\begingroup$ @StianYttervik I'm just going by what the author says in the first sentence: "...although it is not a salt by definition." If you feel that it can be a salt, then that would be great! If so, perhaps you can propose an edit to adjust the wording? $\endgroup$
    – uhoh
    Commented May 5, 2019 at 10:00
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    $\begingroup$ @MathewMahindaratne Excellent, thank you for that! I've made a formatting change to start a new paragraph to improve the visibility. $\endgroup$
    – uhoh
    Commented May 5, 2019 at 10:24
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    $\begingroup$ Ixnay on the ammonium fluoride calculation. Solutions, especially of ionic compounds, are far from ideal when they are as concentrated as quoted in the edit; therefore the linear dependence of freezing point depression on on concentration does not apply. $\endgroup$ Commented May 5, 2019 at 14:32
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    $\begingroup$ @Oscar Lanzi: I agree. It is an important point. Similar to Beer's law. That's why I stated as "theoretically"! :-) Also, thank you uhoh for the edit. $\endgroup$ Commented May 5, 2019 at 16:24

Magnesium perchlorate is far from unique. In fact, if you're willing to spend a little money at that hardware store you could pick up some calcium chloride, whose eutectic reaches about -50°C, not quite as low as magnesium perchlorate but still good enough to cover much of the temperature range on Mars.

Hydrogen chloride, which becomes ionic upon reacting with water, gives a liquid down below -70°C according to a German reference: "Systemnummer 6 Chlor, Ergänzungsband Teil B – Lieferung 1". Gmelins Handbuch der Anorganischen Chemie. Chemie Weinheim. 1968.

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    $\begingroup$ Hydrogen chloride however has a vapor pressure (normal BP is $-85 ^\circ C$) $\endgroup$
    – Buck Thorn
    Commented May 5, 2019 at 17:14

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