[CHMSTRYLover69]

Hi Forum,
I have a relatively simple lab to complete in my undergrad chem class. Here's the situation:
I need to measure the molarity of HCl in an unknown sample to 2 significant figures (but I know it's somewhere between 0.2M and 0.3M). However, all the glassware and pH meters are missing, meaning no titration.
I was thinking of a sort of "titration" with a solid, and instead of using a liquid indicator, we add solid Magnesium, which reacts well with HCl. We add small amounts of Mg until it no longer reacts with HCl, and because of the high disassociation of HCl, we can use stoichiometry to calculate the moles of HCl in our sample (the only problem is finding the volume of the sample).
Does this proposal hold water (no pun intended)?
Are there maybe more efficient or accurate ways of measuring HCl concentration?
Thanks in advance,
CHMSTRYLover

[Borek]

General idea sounds OK, but it will be very difficult to tell when the reaction stops, as it will slow down to a crawl near the "equivalence point".

Sticking with Mg - what are the products of the reaction?

[Enthalpy]

The density of HCL aqueous solutions must be known well enough to tell the concentration accurately from the mass and volume. The refractive index possibly too.

Or could a bicarbonate powder be a faster alternative to magnesium?

[CHMSTRYLover69]

General idea sounds OK, but it will be very difficult to tell when the reaction stops, as it will slow down to a crawl near the "equivalence point".

Sticking with Mg - what are the products of the reaction?

Because HCl disassociates almost completely in water, I believe the reaction will be 2HCl(aq)+ Mg(s) --> H2(g) + MgCl2 (aq), right? As soon as we get Mg(s) that does not react, we can use stoichiometry to determine the moles of HCl in our solution, right?

[AWK]

Dissociation in this experiment is not a significant problem. Acetic acid dissociates in several percents and also reacts completely with magnesium.

To count anything quantitatively you must be able to take advantage of a good weight. Experience with magnesium is not the best. The mass loss of the solution with added magnesium will be very small. On a scale of 100 g of the 0.2-0.3 M acid solution, it will be about 0.03 g. Evaporation of the magnesium chloride solution does not lead to a hydrate of fixed composition.

The addition of excess NaHCO₃ (about 3 g), as suggested by Enthalpy, to 100 g of the solution will lead to the release of gaseous CO₂ and the loss of mass will be about 1.5 g (although the solution will need to be slightly warmed up to about 40 C to remove carbon dioxide almost completely). Based on the known weight of the acid solution, the sample of sodium bicarbonate and the final solution weight, I do hope, you will be able to calculate the HCl percentage and, if necessary, convert it to a molar concentration based on the knowledge of the solution density obtained by interpolation of the table data.

You do not need to weigh exactly 100 acid and 3 g sodium bicarbonate but you need to know the exact masses of the sample weights.