Step 1

**mass = $\frac{\mathbf{4}\overline{)\mathbf{}\mathbf{atoms}}}{\mathbf{1}\mathbf{}\mathbf{unit}\mathbf{}\mathbf{cell}}\mathbf{\times}\frac{\mathbf{1}\mathbf{}\overline{)\mathbf{mol}}}{\mathbf{6}\mathbf{.}\mathbf{022}\mathbf{\times}{\mathbf{10}}^{\mathbf{23}}\mathbf{}\overline{)\mathbf{atoms}}}\mathbf{\times}\frac{\mathbf{102}\mathbf{.}\mathbf{91}\mathbf{}\mathbf{g}}{\mathbf{1}\mathbf{}\overline{)\mathbf{mol}}}$ **

**mass = 6.8356x10 ^{-22 }g**

Step 2

$\overline{)\mathbf{volume}\mathbf{}\mathbf{=}\mathbf{}{\mathbf{a}}^{\mathbf{3}}}\phantom{\rule{0ex}{0ex}}\mathbf{volume}\mathbf{}\mathbf{=}\mathbf{}{(2\sqrt{2}\xb7r)}^{\mathbf{3}}\phantom{\rule{0ex}{0ex}}\mathbf{volume}\mathbf{}\mathbf{=}{(2\sqrt{2}\times 135\mathrm{pm}\times \frac{{10}^{-12}m}{1\mathrm{pm}}\times \frac{1\mathrm{cm}}{{10}^{-2}m})}^{\mathbf{3}}$

**volume = 5.5672x10 ^{-23} cm**

Rhodium crystallizes in a face-centered cubic unit cell. The radius of a rhodium atom is 135 pm.

Determine the density of rhodium in g/cm^{3}.

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