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Name: Ellie F.
Status: student
Age: 12
Location: N/A
Country: N/A
Date: 3/13/2004

For my science project I decided to measure the effect of 0.1, 0.5, and 1M concentrations of NaCl, Borax, table sugar, and NaHCO3 on boiling point. I got a thermometer that is accurate to 0.1 degree C, measured the boiling point of pure water, then measured each of the solutions in the same vessel. Finally I measured the boiling point of water after the experiment and it was the same. As we predicted, the boiling point of sugar, borax, and salt increased about 1C for the 1M solution. But low and behold, the bp for NaHCO3 was depressed by 1.2C for the 1M solution. The solubility of NaHCO3 is about 10g/100 ml water, and we added 8.4g/100ml, so it should have all gone into solution. We repeated the experiment, and the same thing happened.Everything I have read says that adding a soluble substance to water should increase the boiling point. My mother thinks it may be because the NaHCO3 is not dissolving and acts as nucleation points for the water vapor, but it looked like it all dissolved. When it boiled, it did look sort of frothy. What is going on here???


Sodium bicarbonate can break down to produce carbon dioxide gas -- thus, the froth you saw and the apparent lower boiling point.

ProfHoff 826

Your science project has led you into the "unexpected", which is what a good science project should. If all experiments did was to confirm what we already know, science would make no progress.

I believe what is happening is this: NaHCO3 involves several chemical equilibria:

NaHCO3 = Na(+1) + HCO3(-1) dissolved in water
HCO3(-1) + H2O = H2CO3 + OH(-1) dissolved in water
H2CO3 = H2O + CO2 dissolved in water
CO2 (dissolved in water) = CO2 (gas)

The key reaction is the last one. When you heat the solution of sodium hydrogen carbonate to the boiling point, CO2 is driven from the solution into the atmosphere (gases are less soluble as the temperature increases). This increases the concentration of all the other things in solution, thus increasing the boiling point. In addition, a nominal 1 molar solution of sodium hydrogen carbonate also contains dissolved CO2 that does not get driven off by heating. This dissolved CO2 also contributes to the increase in the boiling point. I do not know the values of the pressure of CO2 in equilibrium with the solution, but I am pretty sure that is what's going on.

Vince Calder


Good question. In this case, you have another reaction happening. It just so happens that NaHCO3 decomposes to CO2 when heated in solution. Therefore, as you heat the solution, the solute concentration decreases. This is the source of the frothiness you observed.

Best, Prof. Topper

A broad investigation and a good question, Ellie.

I wonder if some of the carbonate is dissociating, making CO2 gas bubbles and NaOH solution. If so, your water will be really alkaline after boiling, and make broad-range pH paper show pH 12-13. To get it neutral again would then need addition of around 0.01-0.1 mole of any acid, per liter of your bicarbonate solution. Better start slow, 1 drop & swirl, because 0.01 moles of CO2 might have made a big amount of gas bubbles. 5% Vinegar, diluted 10:1, would be fine. If I'm right, the first drops of vinegar you add won't make the usual amount of CO2 bubbles. Continuing to add, after you have added enough to be equivalent to the CO2 you boiled off, more vinegar will make bubbles as usual.

I suspect that NaOH is a little bit "soapy" and encourages the frothiness you described. Make an 0.01 M NaOH solution and shake it to see if that looks true. And of course, emitting CO2 gas in addition to H2O vapor might help make lots of small bubbles, which always looks a bit frothy.

You might want to add NaOH to your list of substances to test for boiling point effect, since the NaHCO3 might be making it.

The CO2 bubbles coming out of solution will then be nucleation points for water vapor. Even more than nucleation points, if the bubbles are about 4% CO2 and 96% H2O, that would reduce the boiling point of the water itself by about 1 degree C. I got that from a "Vapor Pressure Table for Water" in my CRC chemistry reference book.

On the other hand, your mother might have a point, there. 8.4 gm/100ml is pretty close to the solubility limit of NaHCO3. As water boils and flies away, the bicarbonate cannot leave with it. Some water in the immediate area of the bubbles is likely to be supersaturated with bicarbonate, until later when it is stirred with the rest of the bath. Then the bicarbonate would try to un-dissolve, and indeed make nucleation points nearby. It would be a lot like the sea-foam on ocean waves at the beach, frothy. For a substance this close to it's solubility limit, many investigators would omit the case of 1M concentration, and rely instead on 0.5 and 0.25 M.

Serious no-fun scientists would therefore not allow themselves the luxury of telling boiling point by the presence of bubbles! ("Bubbles are too messy.") So they would stir the solution like a tornado. Then there would be no bubbles, just steam evaporating off the fast-swirling surface. And excess bicarbonate on the surface would be re-mixed with the bath very fast. How to tell whether it is boiling? At boiling temperature the vapor pressure is equal to (a maybe a tad higher than) surrounding atmospheric pressure. They would cap the pot with a lid, and set the heat input so a small flow of clear, featureless steam continuously streamed out of a hole in the lid. They would probably even measure how much extra pressure happened inside because of this small hole and flow. But for your class you do not need to do all this. Like I said, too much work and no fun.


Jim Swenson

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