Question

One liter of oxygen gas at standard temperature and pressure has a mass of 1.43 g. The same volume of hydrogen gas under these conditions is 0.089 g. If both volumes contain the same number of gas particles (according to Avogadro's hypothesis), how can this difference in mass be explained?

Answer 1

Indeed, the two samples should contain about the same number of gas particles. However, the molar mass of
`\rm O_2\; (g)` is larger than that of
`\rm H_2\; (g)` (by a factor of about
`16`.) Therefore, the mass of the
`\rm O_2\; (g)` sample is significantly larger than that of the
`\rm H_2\; (g)` sample.

Step-by-step explanation:

The
`\rm O_2\; (g)` and the
`\rm H_2\; (g)` sample here are under the same pressure and temperature, and have the same volume. Indeed, if both gases are ideal, then by Avogadro's Law, the two samples would contain the same number of gas particles (
`\rm O_2\; (g)` and
`\rm H_2\; (g)` molecules, respectively.) That is:

`n(\mathrm{O_2}) = n(\mathrm{H}_2)`.

Note that the mass of a gas
`m` is different from the number of gas particles
`n` in it. In particular, if all particles in this gas have a molar mass of
`M`, then:

`m = n \cdot M`.

In other words,

• 
  `m(\mathrm{O_2}) = n(\mathrm{O_2}) \cdot M(\mathrm{O_2})`.
• 
  `m(\mathrm{H_2}) = n(\mathrm{H_2}) \cdot M(\mathrm{H_2})`.

The ratio between the mass of the
`\rm O_2\; (g)` and that of the
`\rm H_2\; (g)` sample would be:

`\begin{aligned}& \frac{m(\mathrm{O_2})}{m(\mathrm{H_2})} = \frac{n(\mathrm{O_2})\cdot M(\mathrm{O_2})}{n(\mathrm{H_2})\cdot M(\mathrm{H_2})}\end{aligned}`.

Since
`n(\mathrm{O_2}) = n(\mathrm{H}_2)` by Avogadro's Law:

`\begin{aligned}& \frac{m(\mathrm{O_2})}{m(\mathrm{H_2})} = \frac{n(\mathrm{O_2})\cdot M(\mathrm{O_2})}{n(\mathrm{H_2})\cdot M(\mathrm{H_2})} = \frac{M(\mathrm{O_2})}{M(\mathrm{H_2})}\end{aligned}`.

Look up relative atomic mass data on a modern periodic table:

• 
  `\rm O`:
  `15.999`.
• 
  `\rm H`:
  `1.008`.

Therefore:

• 
  `M(\mathrm{O_2}) = 2 * 15.999 \approx 31.998\; \rm g \cdot mol^(-1)`.
• 
  `M(\mathrm{H_2}) = 2 * 1.008 \approx 2.016\; \rm g \cdot mol^(-1)`.

Verify whether
`\begin{aligned}& \frac{m(\mathrm{O_2})}{m(\mathrm{H_2})}= \frac{M(\mathrm{O_2})}{M(\mathrm{H_2})}\end{aligned}`:

• Left-hand side:
  `\displaystyle \frac{m(\mathrm{O_2})}{m(\mathrm{H_2})}= (1.43\; \rm g)/(0.089\; \rm g) \approx 16.1`.
• Right-hand side:
  `\displaystyle \frac{M(\mathrm{O_2})}{M(\mathrm{H_2})}= (31.998\; \rm g \cdot mol^(-1))/(2.016\; \rm g \cdot mol^(-1)) \approx 15.9`.

Note that the mass of the
`\rm H_2\; (g)` sample comes with only two significant figures. The two sides of this equations would indeed be equal if both values are rounded to two significant figures.