Two platinum electrodes are connected to a direct-current power supply and are immersed in a solution of $1.0 \textrm{ M Sn(NO}_3\textrm{)}_2(aq)$ at $25˚C$. As the electrolytic cell runs, tin metal is deposited onto one of the $\textrm{Pt}$ electrodes, and $\textrm{O}_2(g)$ is formed at the other electrode. The two reduction half-reactions for the overall reaction that occurs in the cell are shown in the table below.

$$\textrm{O}_2(g) + 4 \textrm{H}^+ (aq) + 4 e^- \rightarrow 2 \textrm{H}_2\textrm{O}(l) \qquad E°_{cell}=+1.23 \textrm{ V}$$

$$\textrm{Sn}^{2+}(aq) + 2 e^- \rightarrow \textrm{Sn}(s) \qquad E°_{cell}=-0.126 \textrm{ V}$$

A current of $0.850 \textrm{ A}$ passes through the cell for $1.20 \textrm{ hours}$. Assume $100\%$ efficiency.

Calculate the mass, in grams, of the $\textrm{Sn}$ that is deposited on the electrode. In addition, calculate the volume, in liters, of the oxygen gas produced given the temperature is $25.0˚C$ and pressure is $1.00 \textrm{ atm}$.