# Thermodynamics - Is this a correct explanation of the boiling point elevation of salt water?

I am writing a high school assessment on the relationship between salinity of brine and boiling point of the solution, but I am struggling with writing the background, as my teachers are not very familiar with the theory behind this topic.

This is my understanding of why boiling point increases as salinity increases (from what I've gathered on other posts + Google):

Boiling point is defined as the temperature at which the vapor pressure is equivalent to atmospheric pressure. Because the atmosphere essentially exerts a force down on the surface of the water (due to gravitational force pulling down the air), this prevents gas from escaping and the water to arrive to a rolling boil.

By adding salt, the vapor pressure of the resulting solution is lowered, because salt (as an ionic compound) has very low volatility, therefore unlikely to vaporize. As such, the particles must reach a higher temperature (higher travelling kinetic energy) to create the same pressure/force to cancel out the atmospheric pressure.

Set up used:

I made solutions of 5, 10, 15, 20, 25 and 30% salinity. For example: the 5% solution was made with 12.50g of salt dissolved within 250ml of distilled water using a volumetric flask. 50ml of solution was heated up using a Bunsen burner, and the boiling point was recorded as the maximum temperature within 15s of reaching rolling boil. My average boiling point for 5% (50g/dm3) was 102.6 degrees celsius (over 5 trials).

I would greatly appreciate any corrections, or a better explanation of what is occurring. If you could suggest some resources to look at, that'd be awesome.

because salt (as an ionic compound) has very low volatility

is not exactly right.

Table salt in its pure form at room temperature is a crystalline solid owing to its ionic character and is much less volatile than a water (liquid), that by itself need not mean that when you add this to water, it should share this property with water. To be more precise there are a handful of ionic salts that isn't even soluble in water, implying that they don't share their low volatility nature with water.

A more accurate explanation would be in terms of

1. Types of bonding involved (for ionic solutes only): Water molecules in its liquid state is bound by Hydrogen bonding. But when you add soluble salts, because these are ionic (localized non zero charge) it introduces a new kind of bonding called Ion-Dipole Bonds. To cut the story short, since this is also a stabilizing / attracting bond (in fact ion-dipole bond is stronger than Hydrogen Bond) that adds to already present Hydrogen Bonding, it essentially leads to a stronger attractive force that holds water molecules together and so making it harder for them to let go and vaporize. And hence the reason why Brine has lower volatility.
2. Surface area exposed to atmosphere: Here the essential idea is that, evaporation happens only at the surface. Also, solute molecules being more stable in soluble forms donot tend to vaporize when put in a solvent. Keeping these in mind, adding solute decrease the surface area of solvent exposed to atmosphere (or in general a lower pressure region) and so now there are lesser number of solvent molecules evaporating at a time, hence lower volatility.

Before I say anything, I should clarify that I am nowhere near a chemist, its been quite a long time since I handled any sort of equipments. So I have no idea of the precautions necessary/error associated with the various instruments.

Assuming that everything else is foolproof and the physics part is the only missing link - all I can say is about your temperature measurement part. If you measured the temperature 102.6 while the Bunsen burner is burning, its likely that you will get a wrong higher reading.

The idea is this - you have a heat source running at one end and a heat sink at the other end. Now what you are essentially measuring with a thermometer is the Temperature of the solution at Thermal equilibrium and NOT the BP. The reason for this is that even though the liquid solution can only attain its BP temperature, the area close the the burner would be steaming, now steam can attain a lot more temperature (maybe even close to temperature of the burner). So if you measure the temperature while keeping the burner running, what you are effectively measuring is sort of the average temperature of water and the steam (since the burner is continuously producing steam) and naturally its would be higher.

What I suggest you do is to give enough time for the steam to get out, which can only happen if you turn off the burner. Now you wouldn't have to worry about the water getting cooled down in this interval since water have way to high thermal heat capacity that there wouldn't be significant change during this interval.

Now assuming this is the case, I have a fairly good idea on how I would overcome it. I could post it here, but since you are doing a research, I don't want to ruin your experience, so I would highly appreciate it if you come up with a solution to this. Does that sound fair? I suggest that you come up with an idea, edit your original OP and post it there and then maybe tag me in a comment or something.

• How valid is the statement: "you would need to add 230 grams of table salt to a liter of water just to raise the boiling point by 2°C" (ThoughtCo)? Experimentally, I've found that only 50.00g of salt (within 1L) was needed to raise the boiling point by about 2.6°C. It was conducted at atmospheric conditions close to 1atm, so I don't understand where this large disparity is coming from. I also used the boiling point elevation formula to calculate theoretical values, and between 200g/dm3 to 250g/dm3, the elevation was 4.4 - 5.9°C. This is clearly different to the statement, so what's gone wrong?
– Katy
Commented Jun 16, 2021 at 8:42
• @Katy Could you add more details (step by step procedure and the equipments used) on how you performed you experiment. Perhaps you should add that in your original question. Here are some points you should definitely included within your procedure: Did you use pure distilled water? Did you boil around 1L of water? And how... on a stove? How big of a surface was open to atmosphere? How did you measure the temperature and pressure? Try repeating your experiment at least once or twice. Commented Jun 16, 2021 at 16:05
• I've added more information in my original post. I'm very worried because of the large difference between the experimental and theoretical value (calculated using ΔTb = iKbm).
– Katy
Commented Jun 17, 2021 at 2:32
• Great.. I thought this was a home experiment. Anyway @Katy I have updated my answer including what I think might be a possible reason, I wouldn't post it in comment because it was way too long. Commented Jun 17, 2021 at 5:31
• @Katy did you figure it out? Commented Jun 24, 2021 at 16:45

Ions like sodium in water solution have a slight tendency to "hold onto" water molecules. This weak bonding makes it more difficult for the water molecules to escape by evaporation, thereby slightly lowering their vapor pressure.

This effect is called a colligative property of water solutions and other colligative effects include raising the boiling point and lowering the freezing point of water.