at. A discrepancy in this experiment is the one molar solution of strong acid and base in Table 5. The ∆H was C85.0 kJ/mol, more than the three and six molar solutions. This number is wrong. The ∆H for the reaction of the one molar solution of strong acid and base should be below that of the three and six molar solutions, as concentration does affect the amount of heat released by the reaction. This errant temperature could be due to a misreading of the thermometer during the actual experiment. This discrepancy might have been solved by conducting more than a single trial of the reaction under each concentration. However, due to time restrictions, our lab instructor gave us permission to conduct only one trial at each concentration and warned us that our results might be a little strange. A general weakness of the acid base reaction experiments is that, again, our team assumed that the density and specific heats of the solutions were the same as pure water. Also, only one run was conducted on each of the molarities. If more trials had been done, our team would have achieved better results. Additionally, the initial temperature of the base was not factored into our equations. There should not have been much difference, as both solutions should have been at room temperature. However, due to these assumptions, our calculations of ∆H for the acid base reactions are slightly inaccurate. As shown in Tables 4 and 5, the weak acids and bases used in the experiment gave off less heat than the stronger acids and bases used. A weakness of the entire study was the assumptions we made in our calculations of ∆H. Our team assumed that the density and specific heats of our solutions were the same as pure water. The density of pure water is 1.00 g/mL; the specific heat of pure water is 4.182 J/g. According to the lab manual, This is not strictly true, especially as your solutions get more concentrated (...
