Answer:
Approximately 
, assuming that this reaction took place under standard temperature and pressure, and that 
behaves like an ideal gas. Also assume that the reaction went to completion.
Explanation:
The first step is to find out: which species is the limiting reactant?
Assume that 
is the limiting reactant. How many moles of would be produced?
Look up the relative atomic mass of 
, 
, and 
on a modern periodic table:
Calculate the formula mass of :
.
Calculate the number of moles of formula units in 
of using its formula mass:
.
In the balanced chemical equation, the ratio between the coefficient of and that of is 
.
In other words, for each mole of formula units consumed, one mole of would be produced.
If is indeed the limiting reactant, all that approximately 
of 
formula would be consumed. That would produce approximately 
of .
On the other hand, assume that 
is the limiting reactant.
Convert the volume of to 
(so as to match the unit of concentration.)
.
Calculate the number of moles of molecules in that 
of this 
.
Notice that in the balanced chemical reaction, the ratio between the coefficient of and that of is 
.
In other words, each mole of molecules consumed would produce only 
of molecules.
Therefore, if is the limiting reactant, that 
of 
molecules would produce only one-half as many (that is, 
) of molecules.
If is the limiting reactant, 
of molecules would be produced. However, if is the limiting reactant, of 
molecules would be produced.
In reality, no more than of molecules would be produced. The reason is that all would have been consumed before was.
After finding the limiting reactant, approximate the volume of the produced.
Assume that this reaction took place under standard temperature and pressure (STP.) Under STP, the volume of one mole of ideal gas molecules would be approximately 
.
If behaves like an ideal gas, the volume of that of molecules would be approximately 
.
.
.
.