This question is describing an spectrophotometry experiment in which the concentration of copper(II) ions has to be determined when 2.50 g of a copper-rich penny is reacted with 50.0 mL of nitric acid. In addition, the absorbance of a sample of the solution turned out to be 0.12 at 635 nm, so that the concentration was estimated as 12.8 g/L.
First of all, we must recall the Beer's law that can be written as follows:
[tex]A=ebC[/tex]
Whereas A is the absorbance of the sample, e the molar absorptivity, b the cuvette's path and C the concentration. However, this problem is not providing sufficient information to calculate the required concentration, however, similar absorbance experiments have graphed (absorbance vs concentration) calibration curves exhibiting trendlines as the one shown below:
[tex]y=0.009647x-0.003655[/tex]
Thus, since y stands for the absorbance and x for the concentration but in g/L not M (mol/L) for those experiments, we can set y equal to 0.12 so we solve for x as the concentration of the copper(II) ions in the sample:
[tex]y=0.009647x-0.003655\\\\x=\frac{0.12+0.003655}{0.009647}\\\\x=12.8g/L[/tex]
Nevertheless, it would be great if you come up with your own trendline, having clear the units for the absorbance vs concentration graph and thus double check your results because the data was not enough for the time being.
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