The Gibbs energy for the decomposition of $Al_2O_3$ at $500^oC$ is as follows:
$\frac{2}{3} Al_2O_3 \rightarrow \frac{4}{3} Al + O_2$
$\Delta_rG = +960 \ kJ \ mol^{-1}$
The potential difference needed for the electrolytic reduction of aluminium oxide $(Al_2O_3)$ at $500^oC$ is at least ........ $V$.

Based on the following information,arrange four metals,$A$,$B$,$C$,and $D$,in order of increasing ability to act as reducing agents:
$(I)$ Only $C$ reacts with $0.10 \ M \ HCl$ to give $H_{2(g)}$.
$(II)$ When $A$ is added to a solution of the other metal ions,metallic $D$ is formed,but not $B$ or $C$.

The following reaction occurs in a button cell:
$(i)$ $Ag_{2}O + H_{2}O + 2e^{-} \rightarrow 2Ag + 2OH^{-}$
$(ii)$ $Zn \rightarrow Zn^{2+} + 2e^{-}$
Calculate the standard Gibbs free energy change $\Delta G^{o}$ for the overall cell reaction.
$[E^{o}_{Zn^{2+}\mid Zn} = -0.76 \ V, E^{o}_{Ag_{2}O\mid Ag} = 0.34 \ V]$

An aqueous solution of hydrazine $(N_2H_4)$ is electrochemically oxidized by $O_2$,thereby releasing chemical energy in the form of electrical energy. One of the products generated from the electrochemical reaction is $N_{2(g)}$.
Choose the correct statement$(s)$ about the above process:
$A$. $OH^{-}$ ions react with $N_2H_4$ at the anode to form $N_{2(g)}$ and water,releasing $4$ electrons to the anode.
$B$. At the cathode,$N_2H_4$ breaks to $N_{2(g)}$ and nascent hydrogen released at the electrode reacts with oxygen to form water.
$C$. At the cathode,molecular oxygen gets converted to $OH^{-}$.
$D$. Oxides of nitrogen are major by-products of the electrochemical process.

Unlike dry cell,the mercury cell has a constant cell potential throughout its useful life. Why?

Which metal oxide decomposes upon heating?