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The purpose of the following reactions is to make halogen gas:

\begin{align} \ce{2 NaCl (s) + MnO2 (s) + \boxed{\ce{2 H2SO4 (l)}} &-> Na2SO4 (s) + MnSO4 (s) + 2 H2O + Cl2 (g)} \\ \ce{10 KI (s) + 2 KMnO4 (s) + \boxed{\ce{8 H2SO4 (l)}} &-> 6 K2SO4 (s) + 2 MnSO4 (s) + 8 H2O + 5 I2 (g)} \\ \ce{6 KBr (s) + K2Cr2O7 (s) + \boxed{\ce{7 H2SO4 (l)}} &-> 4 K2SO4 (s) + Cr2(SO4)3 (s) + 7 H2O + 3 Br2 (g)} \end{align}

And so we need to use a strong oxidizing agent (in these examples: $\ce{MnO2}$, $\ce{KMnO4}$, $\ce{K2Cr2O7}$) to oxidize halide ions out of their compound.

What I don't understand is the reason why we would need to use sulfuric acid in the reaction if there's another oxidizing agent available and why sulfuric acid is commonly used in this specific redox reaction. Can we use another strong acid instead?

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    $\begingroup$ Yes, it's just acid here. $\endgroup$ – Mithoron Oct 21 '17 at 15:39

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