Evaporation of Organic Solvents

IntroductionWe conducted an experiment to investigate the temperature change associated with the evaporation of several organic solvents and how it was related to strength of the intermolecular forces of attraction in the solvents. We tested a variety of organic solvents (alcohols and alkanes).I initially predicted that pentane would be the most volatile and that 1-butanol would be the least volatile.DataThe data we collected is shown in the table below, along with their respective molecular weights and whether or not the molecule contains hydrogen bonds. More generally, the molecules with hydrogen bonds and higher molecular masses

Substancet1(C1)t2(C1)t = (t1 - t2)(C2)Molecular mass(g mol-1)Hydrogen bonds(Y/N)

ethanol25151046.08Yes

1-propanol2518760.11Yes

1-butanol2622474.14Yes

n-pentane2581772.17No

methanol26101632.05Yes

n-hexane24101486.20No

Analysis

The graph above shows the corresponding temperature change associated with the evaporation of different organic solvents. Pentane has the longest bar and butanol has the shortest; this corresponds to a larger decrease in temperature. The solvents in order of greatest to least temperature change are pentane, methanol, hexane, ethanol, propanol, and butanol.The data also shows that the alcohols generally caused a much smaller temperature change compared to that of the alkenes. This could be because alcohols contain hydrogen and oxygen atoms, which creates hydrogen bonds. Hydrogen bonds are the strongest form of intermolecular forces of attraction, so the alcohols are generally less volatile than the alkenes.

This graph shows the temperature change associated with the evaporation of solvents as well as the molecular mass of each solvent. The varying heights of the pairs of bars on each solvent imply that there is little correlation between molecular mass and temperature change however, our knowledge of hydrogen bonds allows us to discount this seeming lack of correlation. EvaluationOne of the greatest challenges in conducting this experiment was determining the base (starting) temperature of the probe. We decided using only one probe would take too long, as wed have to wait for the probe to return to room temperature after testing each solvent. We alternated between two instead; wed leave one to heat back up to room temperature while the other one gathered data. However, we were not sure when the temperature probe could be used again, as the temperature probes stabilized at different readings even when left in the same environment for a period of time. This systematic error would have resulted in a zero error in the accuracy, affecting each of our readings by the same amount. It may have been helpful to record the temperature of the room using a correctly calibrated instrument and taking readings from the inaccurate instruments in light of the accurate room temperature.Another weakness in the methodology was that our environment was not as controlled as we could have been. When left resting, the temperature probes reading often fluctuated by as much as two degrees Celsius. This may have been due to the open windows nearby, which allowed cold breezes into the room, which may have sped up the rate of evaporation or decreased the temperature of the probe. We could have closed the windows or placed our apparatus in a wind shield of some sort in order to minimize random error. A repeat of the experiment with each different solvent would also serve to minimize random error.