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Balance the following oxidation-reduction equations. The reactions occur in acidic solution. a. $\mathrm{Cr}_2 \mathrm{O}_7{ }^{2-}+\mathrm{C}_2 \mathrm{O}_4{ }^{2-} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{CO}_2$ b. $\mathrm{Cu}+\mathrm{NO}_3{ }^{-} \longrightarrow \mathrm{Cu}^{2+}+\mathrm{NO}$ c. $\mathrm{MnO}_2+\mathrm{HNO}_2 \longrightarrow \mathrm{Mn}^{2+}+\mathrm{NO}_3$ d. $\mathrm{PbO}_2+\mathrm{Mn}^{2+}+\mathrm{SO}_4{ }^{2-} \longrightarrow \mathrm{PbSO}_4+\mathrm{MnO}_4{ }^{-}$ e. $\mathrm{HNO}_2+\mathrm{Cr}_2 \mathrm{O}_7{ }^{2-} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{NO}_3^{-}$

Solution

VerifiedTo balance the oxidation-reduction equations, we have to assign the O. N. to each element and determine the oxidation and reduction. After we are done with that, we have to balance the elements in the equations with H$_2$O and H$^+$ for acidic solution. The last step is to balance the electrons and sum the O and R.

Oxidation is the loss of electrons during a reaction and reduction is the gain of electrons during the reaction.

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