November 22nd - Another Lab! Hooray!

Today we completed an all new lab! This lab was all about molar volume, which is the volume of one mole of a substance. The gas we used was butane, an alkane with the formula C₄H₁₀. Behold:

 

Calculating the molar mass of butane is quite simple. Like we learned in previous lessons, we add up the total molar masses of all the atoms in the compound, like so:

4(12.0) + 10(1.0) = 58.0 g/mol

Starting the lab, we gathered our materials: a butane lighter, water, a graduated cylinder, and a scale. We measured the original mass of the lighter, then displaced the water in the submerged graduated cylinder with butane from the lighter. Due to its gaseous state, it floated to the surface. We measured until it got to about 53 mL of gas, then dried off the lighter and weighed it again. Using that information, we found out the total mass of the butane lighter after the experiment. It ended up being a change 0.14 g. Knowing this, we could calculate the number of moles of butane we used:

0.14 g     x     1 mol     =    0.0024 mol

53.0 g/mol         

Using this information, we calculated the molar volume of butane. However, since the units for molar volume are in L/mol, we must convert the 53 mL of butane we measured into litres and then divide it by our number of moles:

0.053 L   =   22 L

.0024 mol               

 We can then calculate our percent error for the experiment:

[(measured   –   accepted)   /   accepted]   x   100%

[(22L   –   22.4L)   /   22.4L]   x   100%

=   1.8%

There were a few factors, however, that affected our result. For example, we may not have adequately dried off the lighter before measuring for the second time. Another factor is the difference in temperature from STP, which is 0 degrees Celsius at 101.3 kPa.

Wasn't that a fun lab? Here's a video demonstrating the flammability of butane (just because):

 

KA-BOOM!

Posted by Michael.