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Recrystallization Technique

Purifying a solid is an essential skill in chemistry, as well as organic chemistry.  As you will see in future experiments, a product from a reaction is often not pure; sometimes there are byproducts also present, or sometimes there are excess reactants in the reaction mixture.  Thus, we need to be able to separate the product (the compound of interest) from the other substances present.  Recrystallization therefore is a type of separation, the separation of the substance of interest from impurities, which does include any solvents.  In Organic Chemistry labs, the solvent is a liquid which dissolves the solute.  A reminder from General Chemistry — the mixture of the solvent(s) and solute(s) is called a solution.

When recrystallizing a solid, the solvent should be selected such that the compound of interest is essentially insoluble at cold temperatures, but is soluble at higher temperatures.  The impurities should either be insoluble in the solvent at all temperatures, or soluble at all temperatures.  Lastly, the solvent should be easily removable from the solid, which means that it should be volatile (easily evaporates at room temperature). Volatile solvents generally have a low boiling point.  Additional factors in solvent selection, if all else is equal, is to select one that is safe to use and minimizes environmental impact. 

The aspirin solid that is provided to you is not pure; it contains 10% impurities: sugar (sucrose) added to this drug to make it more palatable, as well as a tiny amount of food dye to make it more visually interesting.  It is your task in this experiment to purify the aspirin by removing the sucrose and the dye. Both aspirin and sucrose are white crystalline solids, and so any remaining color in your product is due to residual dye.

Acetylsalicylic acid, or aspirin, is relatively nonpolar, and so it is only slightly soluble in water at room temperature (0.003 g / 1 mL at 25°C).  Its solubility increases at elevated temperatures (for example, 0.010 g / 1 mL at 37°C), according to the literature.  Sucrose is much more soluble in water (2.1 g / 1 mL at 25 °C). The food dyes are highly soluble in water.  Thus, Part 2 of this experiment will allow for a separation between aspirin and the impurities on the basis of their solubility in water using a vacuum filtration process.  Also, most (but not all) of the water used to dissolve the impurities will pass through into the filtrate.  

Water has a relatively high boiling point, compared to most solvents, and very low volatility.  It cannot be easily removed from the recrystallized solid, and so Part 3 of this experiment is focused on removing water from it.  Ethanol, as well as other organic solvents, has a lower boiling point and is much more volatile, but aspirin is somewhat soluble in it.  So, we will use a small amount of ethanol to aid in the removal of the solvent without much loss due to dissolution.  Adding the ethanol slowly (in increments) is a good technique to add just enough solvent to dissolve the aspirin, and avoid adding too much.  If too much ethanol is added and the aspirin does not re-precipitate, do not worry, as it can be removed by allowing some to boil off.  The volatility of ethanol is why we use a watch glass to cover the beaker in Part 3.

Additional Techniques: Filtration and Melting Points 

After the solid has dissolved, we will cool it down slowly to improve the crystallization process, as well as to avoid any safety issues by putting an extremely hot flask into an ice-water bath.  Water is needed in an ice-water bath to improve the amount of contact area between the ice and the flask.  So, allow for some space for water, but not too much otherwise your flask may tip over.

The goal in the recrystallization is to form large crystals, as fine, powdery crystals easily get lost in filtration.  Additionally, larger crystals indicate better purity, as impurities “break” the crystalline structure.  Aspirin crystals have a characteristic needle-like shape.  

To physically separate the crystallized solid from the solvent, vacuum filtration is employed.  This simple technique uses a piece of filter paper to prevent any small solid particles from passing through the holes in the Büchner funnel.  The filter paper should be dampened before the filtration with the solvent to form a good seal between the filter paper and the funnel.  The vacuum needs to be started before the solution is poured into the funnel to maintain that seal between the filter paper and the funnel (the filter paper could get dislodged!).  Often, the solvent is volatile, so start the vacuum just before pouring the solution in. Otherwise if the vacuum runs too long, the solvent in the filter paper could evaporate and the seal would be lost. 

An additional benefit of vacuum filtration vs. gravity filtration (besides the faster speed of filtration) is that allowing air to flow through the solid will help to dry both the filter paper and the solid faster.  If water is the solvent, the paper will take significantly longer to dry, thus a volatile solvent is typically used for recrystallizations. (Gravity filtrations will be introduced in Exp. 4.)

In this experiment, the purity of a solid will be determined by comparing its experimental melting point to the melting point found in the literature.  The melting point (which is sometimes given as a range) is essentially a measurement of how much energy it takes to break the intermolecular forces (IMF) between the molecules in the solid.  This value is unique to every compound, and varies depending upon many factors: the size of the compound, the strength of the IMF, and the way that the compound “stacks” in its solid state.  An impurity (which is any other compound) generally lowers the melting point by interrupting the IMF present in its perfect crystalline form, which results in there being a lower amount of energy required to melt the solid.  (An exception to this is if the impurity is present in a large quantity and has a significantly higher melting point.) Refer to the following figure for the stacking in an example perfect crystal (left) and a crystal with an impurity (right).

Learning Goals

  • Learn how to set up organic lab glassware for a recrystallization and learn to use clamps.

  • Understand the concepts behind recrystallization (e.g., separations based upon solubility in terms of temperature and polarity).

  • Perform a recrystallization.

  • Calculate percent recovery.

  • Pack a melting point tube and measure the melting point of a solid.

  • Understand how the melting point of a solid is related to its purity.