Question

You are given a 2.180 g sample containing a mixture of XO and X203. Ît takes 0.0150 mole of K2Cr207 to oxidize the sample completely to form XO, and Cr*. If0.0187 mole of XO is formed, what is the molar mass of X? (See next page for some background information and help.) Partially balanced Net Ionic Equations for the simultaneous reactions sum of all coefficients = 1 12; total charge on each side = +24 sum of all coefficients-86; total charge on each side +18

Answer 1

Mass of sample (xo & x_2o_3) = 2.180 g moles of k_2cr_2o_7 = 0.0150 mol moles of xo_4^- formed = 0.0187 mol Balancing the reaction - xo(s) + cr_2o_7^2 - + 34 h_3o^+ rightarrow xo_4^- + cr_3 + + 51h_2o (1) Breaking the reaction x_0 rightarrow xo_4^- cr_2o_7^2- rightarrow cr^3 + Balancing the central atom x_0 rightarrow xo_4^- cr_2o_7^2 - rightarrow 2cr^3 + Balance oxygen (o) by adding H_2o 3H_2o + xo rightarrow xo_4^- cr_2o_7^2- rightarrow 2cr^3 + + 7H_2o \r\n Balance H by adding H^+ 3H_2o + xo rightarrow xo_4^- + 6h^+ 14 h^+ + xo rightarrow xo_4^- + 6h^+ 14h^+ + cr_2o_7^2- rightarrow 2cr^3 + + 7h2o Balance change by adding e^- 3h_2o + xo rightarrow xo_4^- + 6h^+ + 5e^- } times 6 6e^- + 14 h^+ + cr_2o_7^2 - rightarrow cr_2o_7^2 - rightarrow 2cr^3 + + 4h_2 o } times 5 6xo + 5cr_2o_7^2- + 34 h^+ rightarrow6xo_4^- + 10 cr^3 + + 17 h_2 o 6xo + 5cr_2o_7^2- +