A typical lab experiment involves the addition of ammonia in excess to solutions of copper(II) to produce a complex ion whose absorbance at roughly 610 nm varies by the concentration of copper in solution.


Why is it less practical to simply test the aqueous solution of $Cu^{2+}$ rather than complexing it first with ammonia?


My thoughts on this begin with the Beer-Lambert law: $$A =\epsilon lc$$

If the molar absorptivity of the ammonia complex in the visible spectrum around 600 nm is greater than that of the aqueous complex, [1] then small changes in concentration of copper can be differentiated by large changes in absorbance. This increases the sensitivity of spectrophotometric methods of determining copper(II) concentrations.


[1] http://chemed.chem.purdue.edu/demos/main_pages/18.8.html

  • $\begingroup$ If you have data that indicate the absorptivity is higher, you answer is correct. But you might be missing the higher solubilization as a small component to this as well. $\endgroup$
    – Lighthart
    Aug 14 '15 at 22:02

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