Why Does ice Melt Faster in tap Water Than in Salt Water?
Most recent answer: 10/22/2007
- Greg Baxley (age 31)
Bakersfield College, CA
Thank you for your correction! We have removed the original answer from our database.
It is true that an ice cube will melt much faster in tap water than in salt water. And although there is a difference in how high the ice cube will float in each (as I’d said before), it is not enough to make the major difference.
When you think of a regular ice cube melting in a regular glass of water, you have to remember that cold water (like the water from the ice cube) is actually denser than warm water (like the water in the cup). This is because in the cold water, the molecules have less energy and are actually closer together than in warmer water. So as the ice cube melts, the cold water coming off of it sinks to the bottom of the glass and the warm water from the bottom comes up to take its place. The water in the glass is actually constantly moving, keeping the ice cube warm by something that scientists call ’convection currents.’
But salt water is much denser than tap water, warm or cold. So when you put a freshwater ice cube in a glass of salt water, the cold water coming off the ice cube doesn’t sink at all. Instead, the dense salt water stays at the bottom of the glass and the cold water stays on the top. Without any convection currents to carry the cold water away from the ice cube, the ice cube melts much more slowly.
Thank you again for this correction... please feel free to let us know if you see any other errors on our page.
If you’re interested in finding more information on experiments to use in your class on the role of convection in melting ice, you can also check out this article: ’Denison, R.F. Agricultural physics: a laboratory mystery to learn about convection. Golden Slate (Calif. Agric. Teachers’ Assoc. newsletter), March, 1999.’ (Submitted by Ford Denison, Age 47 from UC Davis.)
(published on 10/22/2007)
Follow-Up #1: salt and melting ice
- Marcus (age 41)
By the way, the main reason for the lowering of the freezing point by solutes is the effect of the changing liquid volume on the entropy of the solute. This effect is rather easy to calculate, and turns out to agree well with the observed effects. Other effects can matter at higher solute concentrations, and can either increase or decrease the lowering of the freezing point.
The second point concerns the rate at which ice will melt to liquid in a glass of salty or pure water. Rates are much more complicated things than thermodynamic equilibria, and it turns out that in this case the actual effect is typically as described in our old answer.
(published on 03/27/2008)
Follow-Up #2: salt water freezing again
- Binky (age 15)
First, I assume that by "water and ice" you meant "water and salt", otherwise it doesn't connect with the rest of your discussion.
Let's start with the obvious parts based on familiar facts.
If salt tended to stabilize ice more than water, by the fancy mechanism you describe, adding salt to ice would help keep it frozen. Somebody better tell the Highway Department, because they've been using salt to melt ice all these years.
Working our way toward slightly less familiar facts, the mechanism you describe would only work if the salt was actually in the ice, helping line up those water molecules. However, virtually none of the salt goes in to the ice. You can check this by partially freezing some salt water. You'll find that the non-frozen liquid is saltier than what you started with, because salt was excluded from the ice.
I'm trying to see if there is any grain of truth in what you've written. It is true that dissolving table salt in water will lower the temperature of the water, because it takes energy (more precisely, enthalpy) to pull those ions apart. That could help freeze the water. However, that effect amounts to only about 0.9 K cooling for 1 M salt concentration. (That cooling is temporary and has no effect on the freezing time for the salt water if you let it stand at room temperature before putting it in the freezer.) The same 1M NaCl lowers the freezing point of water by about 3.7 K, so it still needs to cool more than plain water before it freezes, even if you didn't let it reach room temperature first.
So I have to conclude that everything you wrote about the science of water and salt is incorrect.
Is it possible that your glass of salt water froze first? Maybe- freezers don't cool evenly, some spots already have frost, some are in good contact with metal, so it takes a lot of care to check whether the result depended on details of placement, etc. However if you got the water really salty, stirring in NaCl until no more would dissolve, the water wouldn't freeze until it was cooled to -21.1 °C. Ordinary home freezers don't usually get quite that cold, so unless you have a very good freezer on the coldest setting you would never have seen it freeze.
For anybody who wonders who is right: do the experiment yourself.
(published on 09/11/2009)
Follow-Up #3: experiments on melting ice
- Bennett Willis (age 68)
Lake Jackson, TX, USA
I wonder if we can use the technique to demonstrate diffusion in class. Our other diffusion demos are messed up by convection, as you say. The issue will be if we can get the apparatus to sit untouched for a few days between classes.
(published on 09/15/2009)
Follow-Up #4: stirred salt water and ice
- James M (age 60)
(published on 11/05/2009)
Follow-Up #5: ice melting with salt or not
- Jim Houghtaling
Lancaster, SC 29720
(published on 03/04/2010)
Follow-Up #6: cooling ice with salt
Incidentally, adding ordinary table salt (NaCl) to liquid water soaks up some enthalpy, so if the salt and water started at the same temperature, they get cooled a bit by forming a solution. Some other salts release some heat as they dissolve.
(published on 05/16/2013)
Follow-Up #7: disputes about melting of ice by salt
- Billy (age 21)
In looking over your discussion, there seems to be one point on which we clearly disagree. You write "One major assumption that is incorrect in seemingly all of these answers is that the melting point and freezing point are the same. " The equality of the thermodynamic melting and freezing points is not just a feature of some detailed model. It follows directly from the fundamental laws of thermodynamics. For a given solution concentration one phase or the other has the lower free-energy at each temperature, and that phase is the stable one. There's just one temperature where the free-energies are equal, and that's the melting/freezing point. For non-saturated solutions, the concentration changes as the freezing or melting occurs, giving a melting/freezing range. However, at each temperature in that range a precise fraction of the material is melted in equilibrium, regardless of previous history.
Interactions between ions of salts do change the magnitude of the freezing point depression and boiling point elevation, giving a correction to the simple colligative values except for very dilute solutions. The type of argument you give, trying to pick up various terms in the dynamic behavior separately, is tempting to use for pictures but famous for leading to wrong conclusions. The deep and rigorous rules of thermodynamics are a reliable guide to equilibrium properties, although not to rates.
(published on 07/27/2013)