A researcher mixed 5.00mL of 3.33x10-1M Fe(NO3)3 with 50.00mL of 4.44x10-3 M KSC

Question

A researcher mixed 5.00mL of 3.33x10-1M Fe(NO3)3 with 50.00mL of 4.44x10-3 M KSCN, using 5.00x10-1M HNO3 as solvent for both solutions. The resulting measured absorbance for the solution was 0.315. Based on the given data, calculate [Fe*]*[SCN*].

Explanation / Answer

The net ionic reaction is (Fe3+) + (SCN-) ===> Fe(SCN)2+ You need to assume that the reaction has gone to completion in tube (i); since there is so much SCN-, the reaction will be driven to the right, using up all of the Fe3+ to produce Fe(SCN)2+. Since the Fe3+ came from the Fe(NO3)3, that's what they mean by [ Fe(SCN)2+] final = [Fe(NO3)3] initial. The initial [Fe(NO3)3] in tube (i) is (1 ml x 0.5 mM)/(12 mL), so the final [Fe(SCN)2+] is equal to that. The initial [Fe(NO3)3] in tube (ii) is (2 ml x 0.5 mM)/(12 mL), but the reaction does not go to completion; instead, an equilibrium is established, with a lower [Fe(SCN)2+] than tube (i) Tubes (ii) through (v) should have been lighter in colour than tube (i), with lower absorptions. The the ratio of the absorptions is equal to the ratio of the equilibrium concentrations of Fe(SCN)2+ (ie. if tube (ii) had half the absorption of tube (i), it would have half the concentration of Fe(SCN)2+ of tube (i).

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