Making Solutions

[Mark_Botros_1D]

I was wondering if someone may share their ideas upon why chemists do not typically measure out very precise, predetermined masses when making solutions. Thank you.

[Mark_Botros_1D]

Also, I believe this was brought up in the Mixtures and Solutions section in the textbook.

[Devika Nair 2D]

"A Note on Good Practice: The laboratory procedure is to take the approximate amount of solute needed, weigh it accurately, then make up the solution, calculating the actual concentration of solute from the mass that was used and the final volume of the solution. For example, you might find that you had measured out 2.403 g, in which case the molar concentration would be 0.0385 M Cu(SO4)(aq)."

Is this the section you were referring to? If so, I think your answer is just practicality. Real-world labs have time and money constraints. Most lab balances measure to 3 or 4 decimal places, and it is incredibly difficult to weigh out solids to an exact thousandth or ten thousandth of a gram. Even if the chemist manages to do so (without wasting potentially expensive material), it is not worth the time spent. For most experiments, only an approximate mass is needed. As long as it isn't an egregious difference from the calculated mass, you'll find that, in the end, you end up with a solution that is really close to the target concentration. In the example the textbook provided, the actual and target concentrations only differed by 0.0005 M, which is negligible.

Save time and approximate.

[Mikayla 2G]

Additionally, it is pretty simple to re-calculate the molarity of the solution and use the new molarity in calculations for the experiment rather than spending a lot of time trying to adjust the sample by thousandths of a gram.