Sciencemadness Discussion Board

Diluting HCL produces heat

herschel - 6-9-2016 at 01:03

If I take 36% HCL and dilute it 50/50 with water, heat is produced. Why?

RocksInHead - 6-9-2016 at 03:18

Because a reaction is taking place. While the H+ and Cl- ions are being dissolved in the water, water is taking some of those H+ ions to form a complex- H3O+, a hydronium ion. This reaction is exothermic, hence why heat is produced when you dilute strong acid with water.

Hope this helps!

[Edited on 6-9-2016 by RocksInHead]

herschel - 6-9-2016 at 12:50

Does this mean that in 36% HCL some of the HCL is dissolved without being dissociated?

vmelkon - 7-9-2016 at 11:52

No, when you dissolve HCl in water, it is 100% dissociated. It is considered a strong acid.
Adding more water probably causes a hydration reaction (Cl- gets surrounded by more H2O).

DraconicAcid - 7-9-2016 at 12:14

Whenever you mix two liquids, there will be some heat absorbed or generated. You are breaking some intermolecular attractions between the original liquids, and forming new ones between the two liquids. These will be of different strengths (unless it's an ideal solution), so some energy will be lost or absorbed.

There doesn't have to be a reaction for this to take place- ethanol and water will have a measurable temperature change when mixed (I know this, because I remember having my students measure it one year).

woelen - 7-9-2016 at 23:40

Quote: Originally posted by herschel  
Does this mean that in 36% HCL some of the HCL is dissolved without being dissociated?
I disagree with vmelkon. In 36% HCl, some of the HCl indeed is not ionized and exists as molecules of HCl. The major part is ionized, but maybe 10% or so still exists as HCl. This also is the reason why 36% HCl fumes heavily. Ionized HCl cannot escape the liquid, only neutral molecules of HCl can escape it.

The same is true for 65% nitric acid. Part of it exists as covalent molecules HNO3. That is why 65% nitric acid is such an energetic oxidizer. Nitrate ions hardly are oxidizing in water, it is the molecule HNO3 which is.

j_sum1 - 8-9-2016 at 02:19

I think it is a lot less than 10%, woelen.

Based on a density of 1.1789 for 36% HCl I calculate a molarity of 11.64M
Given a pKa of -6.3, this works out to be a pH of -1.0659. Or more usefully, the concentration of hydrogen ions is 11.64 moles per litre.

These two 11.64 figures are not identical. There is a difference of some 6.8×10–5. Which effectively means 0.000068 moles of the HCl is undissociated at that high concentration. This is 0.00058% of the HCl present.

Of course it is an equilibrium and so as the HCl gas escapes more molecular HCl is created to replace it.
I suspect that it is the rate of diffusion that is the limiting factor for the escape of HCl from the solution. I stand to be corrected on this. (And I stand to be corrected on my calculations too.)


Metacelsus - 8-9-2016 at 05:07

I'd be wary of calculating the undissociated HCl concentration like that, because the solution will be very highly non-ideal. I wouldn't be surprised if a few percent remained undissociated.

DraconicAcid - 8-9-2016 at 07:29

Quote: Originally posted by j_sum1  
I think it is a lot less than 10%, woelen.

Based on a density of 1.1789 for 36% HCl I calculate a molarity of 11.64M
Given a pKa of -6.3, this works out to be a pH of -1.0659. Or more usefully, the concentration of hydrogen ions is 11.64 moles per litre.

That pKa is for aqueous solution. If you've changed to 36% HCl, that's not aqueous. You've got, what? four molecules of water for every hydronium/choride ion pair? The Ka will be different.

j_sum1 - 8-9-2016 at 08:05

Well, I did invite correction. Thanks.

How would you calculate an acid that concentrated?

woelen - 9-9-2016 at 00:01

Indeed, the concept of pH and the use of pKz only works fine at low concentrations. Up to appr. 1 mol/l the approximations still are acceptable, but at higher concentrations the error quickly grows with concentration.

There is no standard framework for doing this kind of calculations which works for all acids. Each acid has its own peculiarities and modes of non-idealness. Some modes are:
- no splitting of acid molecule in ions
- formation of ion-pairs (electrostatic forces keep two ions of opposite charge close to each other, which makes them unavailable for reactions)
- formation of acid-specific complexes.
- dissociation in anhydride/oxide and water, which makes the acid unavailable (e.g. H2CO3, HNO2, H2SO3)
The higher the concentration, the stronger the effect of these non-ideal behaviors.

For most acids, the non-ideal behavior leads to less acidity than one would expect on the basis of pKz only. An exeption is HF. At high concentrations this acid becomes stronger, a concentrated solution is more acidic than one would expect theoretically.