How could mass increase when sulfuric acid is added to calcium carbonate?

Context: I was doing an experiment in class that investigates the rate of mass loss over time when $$\pu{20 mL}$$ of sulfuric acid is added to $$\pu{5g}$$ of calcium carbonate. I did the experiment 3 times, each experiment involves mixing four, $$\pu{20 mL}$$ different concentration of sulfuric acid ($$\pu{0.5 M}$$, $$\pu{1.0 M}$$, $$\pu{1.5 M}$$, and $$\pu{2.0 M}$$) into $$\pu{5g}$$ of calcium carbonate.

Hypothesis: If concentration increase, the rate of mass loss should increase too

Result: Everything went out just fine, however, in each of the 3 experiments, when I add the $$\pu{2.0 M}$$ sulfuric acid to the calcium carbonate, the mass INCREASES (mass still loss overtime).

Could anybody please explain how could this be happening? As they react, carbon dioxide is produced and there should only be a mass loss. Also, this only occurs with the $$\pu{2.0 M}$$ sulfuric acid, the other concentrations work out just fine.

Could the electronic scale wasnt working properly? but this only happened with $$\pu{2.0 M}$$ acid.

• I'm not sure how you were taking the weights. Did you account for the fact that the densities of the sulfuric acid solutions are different? In other words 20ml of 2 molar acid weighs more than 20 ml of 0.5 molar acid. – MaxW Mar 31 at 17:10

As MaxW pointed out in his comment, I also felt like this may be the case of densities of $$\ce{H2SO4}$$ solutions. Thus, I googled for densities and found few sites listing them. In this Steffen's Chemistry page, you'd find some trends. I just listed relevant densities and milarities of some 1-15% $$\ce{H2SO4}$$ solutions, and graphed them molarity vs density to find the relationship. I used data for 1-25% to graph for better reliability. Sure enough, there was a linear relationship of $$y=0.0597x+1.0006$$ with $$R^2=0.9998$$ reliability:

So, using the graph, I calculated the masses of your $$0.5, 1.0, 1.5, \text{and }\pu{2.0 M}\; \ce{H2SO4}$$ solutions. They were $$20.609, 21.206, 21.803, \text{and }\pu{22.40 g}$$, respectively. Thus, if you are using electronic balance with $$\pu{0.10 g}$$ accuracy, you may have readings of $$20.6, 21.2, 21.8, \text{and }\pu{22.4 g}$$ for one set. All except for last have closed to $$\pu{21 g}$$ weight. Last one is $$\pu{22 g}$$, a little bit more than others. Make sense? :-)

• WOW Cheers Mathew! I am at school right now and i just saw this! You saved my life! – Fred Weasley Apr 1 at 3:33
• Does it matter if they are closer to 21 or 22? – Fred Weasley Apr 1 at 10:21
• I think you are taking mass difference to calculate rate. I think you are okay on this. – Mathew Mahindaratne Apr 1 at 16:41

There are several contributing factors to this outlier observation. First of all, when scientists did mass based experiments, even hundreds of years ago, they recorded mass up to four decimal places. The mass of CaCO3 should have been recorded to 0.1 mg for example 5.0612 g. When you say that mass increases, we do not know the details of mass change. In short, never record mass with one significant figure.

Secondly, how did you measure volumes of 20 mL acid? If you were using cylinders, they are never accurate. There is significant error in volume transfer. With a relatively viscous solution such as 2M H2SO4, if you transfer 20 mL, the transferred volume via cylinder is usually less than 20 mL.

This all about good laboratory practice. Coming to chemistry: when CaCO3 reacts with H2SO4, if forms an insoluble salt called calcium sulfate (CaSO4). With concentrated 2 M sulfuric acid, once it comes in contact with CaCO3, it may form a coating of CaSO4, which would slow down the rate of reaction. Did you see some unreacted solid mass?

Also conc. sulfuric acid is hygroscopic.

• Thanks Mr.Farooq! I will make sure I apply these tips in the next experiment! – Fred Weasley Apr 1 at 11:07