Spectrophotometric determination of copper concentration

Context:
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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.


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


Thoughts:
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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.


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


