Calculate $E_{\mathrm{cell}}^{\circ}$ for each balan ced redox reaction and determine if the reaction is spontaneous as written. a. $\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)+4 \mathrm{Ag}(s) \longrightarrow 4 \mathrm{OH}^{-}(a q)+4 \mathrm{Ag}^{+}(a q)$ b. $\mathrm{Br}_{2}(l)+2 \mathrm{I}^{-}(a q) \longrightarrow 2 \mathrm{Br}^{-}(a q)+\mathrm{I}_{2}(s)$ c. $\mathrm{PbO}_{2}(s)+4 \mathrm{H}^{+}(a q)+\mathrm{Sn}(s) \longrightarrow \mathrm{Pb}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{Sn}^{2+}(a q)$

Calculate $E_{\mathrm{cell}}^{\circ}$ for each balanced redox reaction and determine if the reaction is spontaneous as written. a. $2 \mathrm{Cu}(s)+\mathrm{Mn}^{2+}(a q) \longrightarrow 2 \mathrm{Cu}^{+}(a q)+\mathrm{Mn}(s)$ b. $\mathrm{MnO}_{2}(a q)+4 \mathrm{H}^{+}(a q)+\mathrm{Zn}(s) \longrightarrow \mathrm{Mn}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{Zn}^{2+}(a q)$ c. $\mathrm{Cl}_{2}(g)+2 \mathrm{F}^{-}(a q) \longrightarrow \mathrm{F}_{2}(g)+2 \mathrm{Cl}^{-}(a q)$

Balance each redox reaction occurring in basic aqueous solution. a. $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{Br}^{-}(a q) \longrightarrow \mathrm{MnO}_{2}(s)+\mathrm{BrO}_{3}(a q)$ b. $\mathrm{Ag}(s)+\mathrm{CN}^{-}(a q)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{Ag}(\mathrm{CN})_{2}^{-}(a q)$ c. $\mathrm{NO}_{2}(a q)+\mathrm{Al}(s) \longrightarrow \mathrm{NH}_{3}(g)+\mathrm{AlO}_{2}(a q)$

Determine whether or not each metal dissolves in 1 M $HIO_3$. For those metals that do dissolve, write a balanced redox equation for the reaction that occurs. a. Au b. Cr

Determine whether or not each metal dissolves in 1 M $HNO_3$. For those metals that do dissolve, write a balanced redox reaction showing what happens when the metal dissolves. a. Cu, b. Au.