# With the reaction, determine the masses of HSO_3NH_2 required to produce 0.135 g of N_2(g) collected in a burette above water. How do you calculate the volume that this mass of N_2(g) would occupy at 23.4°C and 7.10*10^2 mmHg barometric pressure?

## $H S {O}_{3} N {H}_{2} \left(a q\right) + N a N {O}_{2} \left(a q\right) \to N a H S {O}_{4} \left(a q\right) + {N}_{2} \left(g\right) + {H}_{2} O \left(l\right)$

Apr 9, 2017

Calculate the masses from the related molar amounts. Calculate the volume of ${N}_{2}$ from the ideal gas laws.

#### Explanation:

The balanced equation shows that every mole of ${N}_{2}$ requires one mole of $H S {O}_{3} N {H}_{2}$. 0.135g of ${N}_{2}$ is $\left(\frac{0.135}{28}\right) = 0.00482$ moles of ${N}_{2}$.

This requires $\left(0.00482 \cdot 97\right) = 0.468 g$ of $H S {O}_{3} N {H}_{2}$ and $\left(0.00482 \cdot 53\right) = 0.42555 g$ of $N a N {O}_{2}$.

The volume of ${N}_{2}$ produced is $V = \left(\frac{n \cdot R \cdot T}{P}\right)$

P = 710/760 = 0.934atm, R = 0.0821 L-atm/K-mol, T = 296.6’K, n = 0.00482 moles, V = Liters

$V = \left(\frac{0.00482 \cdot 0.0821 \cdot 296.6}{0.934}\right)$ ; $V = 0.126 L$