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: 327-2. Related Techniques and Instruments Technique: Small-Scale Crystallization Crystallization is a common method of purif...

: 327-2. Related Techniques and Instruments Technique: Small-Scale Crystallization Crystallization is a common method of purifying organic solids that contain a small amount of impurities. The technique of crystallization takes advantage of the greater solubility of a compound in a hot solvent. A saturated solution at a higher temperature normally contains more solute than the same solute-solvent mixture at a lower temperature. Therefore, when the hot saturated solution cools, the solute crystallizes out of the solution. When the amount of the solid to be crystallized is small (usually < 300 mg), a slightly different version of the crystallization technique is necessary. Even though many steps are the same, the technique appears different because you are working with a much smaller amount of material and different types of smaller containers need to be used. Here are the steps of the small-scale crystallization: 1. Find a suitable crystallization solvent (or solvent pair) as described below. 2. Dissolve the solid in a test tube (or other small-scale container). A test tube rather than an Erlenmeyer flask has been chosen because the size of the glassware should be suited to the amount of solution. In any operation, if you use glassware that is too large, you will loose material unnecessarily over its interior surface.Try to use the minimum amount of hot solvent because less product will be recovered from a solution that is not saturated. 3. Cool the solution undisturbed to room temperature and then ice temperature. Cooling a small amount of solution will probably take no more than 10 minutes. Such cooling periods are convenient for doing other lab work; in this experiment it is suggested that you prepare your wash solvent at this time. 4. Remove the supernatant using an un-chipped pipet: squeeze the air out of the bulb and carefully press the tip of the pipet flat to the bottom of the test tube; slowly release the bulb to remove the supernatant solution by drawing it up into the pipet, keeping the solid in the test tube. Then wash with a small amount of cold solvent. Remove as much supernatant as possible before washing. Divide your wash solvent into several portions, and mix each one well with the solid before you pipet it out. Keep every- thing cold throughout; and if you see an alarming decrease in the amount of product, do not perform additional washes. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.1 Department of Chemistry.  327-2. Related Techniques and Instruments Finding a Proper Crystallization Solvent The most crucial aspect of a crystallization is the choice of solvent. A proper crystallization solvent is one in which the solid compound has a maximum solubility (very soluble) when it is hot and a minimum solubility (insoluble) when it is cold. Table LF.1 lists some common crystallization solvents. Table LF.1: Properties of Common Crystallization Solvents Solvent Formula Boiling Point, °C Safety Considerations Ligroin CnH2n+2 varies (check label) Flammable Hexanes C6 H12 67-69 Flammable Diethyl ether (ether) (CH3 CH2 )2 O 35 Flammable Ethyl acetate CH3 COOCH2 CH3 77 Flammable Methylene chloride CH2 Cl2 40 Flammable, Toxic Acetone (CH3 )2 C – O 56 Flammable Ethanol (95%) CH3 CH2 OH 78 Flammable Methanol CH3 OH 65 Flammable, Toxic Water H2 O 100 – By the principle “like dissolves like”, you can sometimes make a useful prediction based on the polarities of the compounds and the solvents. For example, sucrose (table sugar), a polar compound that contains many - OH groups, dissolves in water, a polar solvent. However, you still must determine empirically how soluble the compound is and whether it can be recovered upon cooling. Finding a useful solvent is a process of systematic trial-and-error. If you find, by the method described below several solvents that appear suitable, you might be better off using one with a moderate boiling point (65-95 °C). If the boiling point is too low, you will have difficulty controlling evaporation during the heating step. If the boiling point is too high, you cannot boil the solution on the steam bath, which is by far the safest and most convenient way of heating. Ideally, a solid to be crystallized should be soluble in a solvent when it is hot (usually boiling) and insoluble in the same solvent when it is cold. In the trial-and-error process of finding a solvent, the idea is to test a small amount of the pure solid in a variety of solvents, both non-polar and polar. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.2 Department of Chemistry.  327-2. Related Techniques and Instruments A good crystallization solvent has two properties: 1) the solid is soluble in the not (boiling) solvent and 2) the solid is insoluble in the cold solvent. Crystallization, as mentioned above, is a slow process and requires patience. Solvent trials, however, can be done quickly. Here are the steps for determining a good solvent: 1. Add a small amount (usually 20-30 mg) of the solid to be crystallized to a test tube. 2. Add approximately 1 mL of a trial solvent. Stopper the tube and shake to mix. If the solid dissolves at room temperature, it is soluble. The solute is actually too soluble in this solvent for a crystallization to be effective. This solvent is not a proper crystallization solvent. 3. If none (or very little) of the solid dissolves at room temperature, unstopper the tube and carefully heat the mixture (do not forget to add a boiling stone or stick) to boiling. If the solid still does not dissolve, then it is too insoluble to be used for crystallization. This solvent is also not a proper crystallization solvent. 4. If the solid does dissolve completely in a hot solvent, plunge the solution into an ice-water bath. The formation of solid suggests that this is a good crystallization solvent. (Of course this solid would be a precipitate, not a crystal, but your purpose here is finding the solvent, not doing a careful crystallization.) If the solid does not form, then this solvent is not a proper crystallization solvent. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.3 Department of Chemistry.  327-2. Related Techniques and Instruments Instruments: How to Use and Read an Electronic Balance Electronic top-loading balances, Figure LF.1, can be read to three decimal places and are widely used in modern organic laboratories. Figure LF.1: Commonly used heating equipment. Follow these general guidelines when using a balance: These balances are expensive and delicate precision instruments. Treat them with care. If anything (solid or liquid) spills on or near the balance, clean it immediately to avoid the corrosion from the spilled chemical. A milligram balance usually has a draft shield. While a sample is being weighed, this shield should be closed to prevent the air draft from disturbing the weighing pan. While recording masses, always record to three decimal places. Be sure to have your ELN with you as you are performing the measurement - do not rely on your memory. No chemical should be weighed directly on the balance pan – always weigh in a glass container (a vial or beaker), in a plastic weighing boat, or on a piece of weighing paper. The mass of the container or weighing paper will be tared (subtracted) by pressing the tare (zero) bar in the front of the balance before the sample is added. If the mass of the container is not tared, it should be determined and recorded separately and subtracted in a later calculation. If a volatile liquid is weighed, a cap or cork for the container should be used so that the sample will not evaporate during the weighing process. The mass of the cap or cork should also be included in the tare. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.4 Department of Chemistry.  327-2. Related Techniques and Instruments Our balances are accurate to three decimal places and you should record your mass to three decimal places. 3.00 g means anywhere between 2.950 g and 3.050 g (recall significant figures and rounding) Example LF.1 1. Always record the exact amount measured in your ELN. Record to three decimal places. 2. Place a glass container or a piece of weighing paper on the balance pan and press the tare bar. The numerical readout will be “0.000” on the digital display. If the balance shows two decimal places, press the tare bar again. 3. Use a spatula to add small portions of sodium chloride until the desired mass is shown on the display. 4. Record the exact mass to three decimal places. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.5 Department of Chemistry.  327-2. Related Techniques and Instruments Instruments: Heating Equipment Some commonly used heating equipment in an organic laboratory are shown in Figure LF.2. Figure LF.2: Commonly used heating equipment. Although it will not be used in this course, the steam bath is the safest heating equipment to use in the organic laboratory. It does not generate an open flame. The heat comes from the steam and only reaches the maximum of about 90-95 °C; therefore, it will not overheat as heating mantles sometimes do. It is very useful for heating low-boiling solvents, especially flammable ones. The heating mantle is one of the most commonly used heating devices in organic laboratory. It is designed to heat the contents of a round-bottom flask. There are two major types of heating mantles available commercially. One type consists of soft fiberglass; each mantle fits a specific size (that is, the heating mantle sized for a 100 ml flask will not fit any other size). The other type has a metal exterior while the inside well can be either hard ceramic or soft fiberglass. The fiberglass heating mantle can be used for multiple sizes, such as the designated and smaller. For convenience and safety, the heating mantle is often supported by an iron ring or a labjack underneath. The electric plug of a heating mantle should always be plugged into a variable transformer (or rheostat), such as a Variac, or other type of controller – never the standard wall socket – to adjust the voltage, and thus, the rate of heating. Hot plates work well for heating flat-bottomed containers like beakers and Erlenmeyer flasks. However, they are not recommended for heating flammable solvents because the very hot surfaces of the hot plates may be a fire hazard. The hot plates used in our laboratory are dual function, they heat and stir solution by using a magnetic stir bar. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.6 Department of Chemistry.  327-2. Related Techniques and Instruments Technique: Melting Point Determination The melting point of a solid is the temperature at which transition from solid to liquid occurs at one atmospheric pressure. However, unlike the effect on boiling point, the effect of a small change in the pressure on melting point is negligible and usually ignored. Each pure compound has its own characteristic melting point. In theory, this value could be measured accurately and reported in the literature. In practice, the measurement is not so accurate; you will often run across slightly different melting point reports for the same compound in different literature sources. The melting point is determined by heating a very small amount of the solid slowly (ideally at the rate of about 1 °C per minute). Although we use the term melting point, what you will actually measure in most cases is a range; that is, the beginning (first appearance of liquid) and end (last disappearance of solid) are not at the same temperature. For instance, a typical melting point range for trimyristin might be observed 53-54 °C. (Note that the term “range” is also used to mean the difference between the high and low value. When the two meanings become confusing, it is convenient to talk about the span of the range, in this example, 1 °C.) Characteristics of Melting Point Each pure compound has its characteristic melting point. The melting point measurement can provide informa- tion about the purity and identity of a compound. A pure organic solid often has a sharp melting point and melts over a range of 1.0 °C or less, while a less pure solid exhibits a broad melting range. Usually, a melting range of 2.0 °C or less indicates a pure compound. An impure compound also exhibits a lower(depressed) melting point compared to that of the pure solid. For example, purified trimyristin might melt at 53-54 °C, but crude trimyristin might have a melting point of 48-52 °C. Melting point information can also be used to identify an unknown solid. There are many reference books and literature papers containing lists of compounds with their melting points. It is not possible to identify an un- known compound without more specific information because there are hundreds of compounds with the same melting point. However, you can determine whether two samples with the same melting point are the same by taking a mixed melting point. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.7 Department of Chemistry.  327-2. Related Techniques and Instruments Determining a Melting Point Preparing the Sample Determination of a melting point is an analytical technique. This means that you need only an extremely small sample to make the measurement. The amount of sample that goes into the melting point tube (or capillary tube) should be of approximately the same height as the width of the tube (~1-2 mm) – you will see that this is on the order of a few milligrams. Figure LF.3 illustrates how to load a melting point tube. Crush a few milligrams of the dry solid into a fine powder, and mix to get a uniform sample. Press the open end of the melting point tube into the mound of powder to make a small plug. Invert the tube and tap it on the bench top so that the sample packs down at the bottom. Then insert the sample tube into the melting-point instrument. There are five holes for tubes, but it is impractical to try to take five melting points at a time. However, you should take two or three at a time whenever you are comparing values (for example the crude and purified trimyristin). Figure LF.3: Preparing a sample for melting point determination Heating the Sample to its Melting Point Turn on the power/light switch. Turn the heat-select knob to raise the temperature quickly to about 10-15 °C below the expected melting point of the sample. Then adjust the knob so that the temperature is rising no more than 1-2 °C/minute. If you are impatient at this step, you may get an artificially high melting point and/or broad range. If the melting point of the sample is unknown, it will save time to take a preliminary melting point first by rapidly heating. (Of course this preliminary melting point will be inaccurate, but it will give you an idea at what Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.8 Department of Chemistry.  327-2. Related Techniques and Instruments temperature, approximately, the sample will melt.) Then take a careful measurement of a second sample. Decomposition at the melting temperature is common especially when impurities are present. For this reason, you should always prepare a fresh sample for each melting point measurement. Recording the Melting Point Record the appearance of the first drop of liquid and the disappearance of the last piece of solid. This is the melting point range. Be sure to record values directly into your notebook, both the beginning and end of melting. Look carefully for the first drop of liquid – shrinking or movement of solid does not count. Always report the melting point as a range even if the observed onset of melting and the offset is the same temperature. An example of the proper way for a student to report an observed melting point range: Melting point (or MP, mp) of crystallized benzoic acid is 121-123 °C. If you are discussing purity, it is often appropriate to cite the span of the range, which is 2 °C in the above example. If your solid happens to melt at one temperature, it should still be recorded with two values. Interpreting the Melting Point If a compound is pure, its melting point value will be close to the one in the literature, and the range will be sharp. How close to the literature value should the melting point be? Because thermometers are usually not calibrated, there is an unknown source of error in any melting point measurement. This error could occur in the literature, in your value, or both. You should allow ± 5 °C as a margin of error (from the offset temperature) at low MP temperatures, and slightly more (about ± 7 °C) at higher temperatures (greater than 150 °C). How sharp should the range be? The sharper the better, although for many purposes as much as 2 °C is acceptable. If the compound is impure, its melting point value will be depressed, or its range will be broadened, or both. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.9 Department of Chemistry.  327-2. Related Techniques and Instruments Example LF.2 Melting Interpretations of Various Situations Your “pure” trimyristin has a melting point of 53.5-54 °C. The literature value is 56 °C. Interpretation: A 2 °C difference from the literature value is insignificant, and the range is 0.5 °C. The product is pure. Your “pure” trimyristin has a melting point of 50-54 °C. The literature value is 56 °C. Interpretation: Although there is only a 2 °C difference from the literature value, the melting point range is 4 °C. Either the product is impure or you took your measurement too quickly. It is worth a check.) Your “pure” trimyristin has a melting point of 52-54 °C. The literature value is 56 °C. Interpretation: It is pure enough for most purposes. You have an unknown whose melting point you have measured at 112-113 °C. Interpretation: The 1 °C range indicates the compound is pure. Therefore, you can have reasonable confidence that the literature value lies ~5 °C from your measurement, that is, between 107 °C and 118 °C. This information will be very useful in identifying the compound. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.10 Department of Chemistry.  327-2. Related Techniques and Instruments Instruments: Melting Point Apparatus Several different types of electrically heated melting point devices are commercially available. The Thomas-Hoover melting point apparatus, (Fig- ure LF.4). In this instrument, capillary tubes con- taining solid samples are submerged in an electrically heated oil bath. A magnifying glass is used to view the crystals in the capillary tubes clearly. In some more advanced versions of the Thomas-Hoover melt- ing point apparatus, a periscope is also provided so that the mercury (or dyed alcohol) thread of the ther- mometer and the crystals in the capillary tubes can be viewed simultaneously. The Barnstead Electrothermal Digital Mel-Temp shown in Figure LF.5 is a popular melting point apparatus model. The digital system features a built-in temper- Figure LF.4: Thomas-Hoover melting point appara- tus ature display, microprocessor-controlled temperature ramping, and rapid fan cooling. Temperature is sensed by a platinum RTD (Resistance Temperature Detector). During a melting point measurement, the on-set (or “start”) temperature and the off-set (or “end”) temperature can be recorded and read from the digital display. Figure LF.5: Barnstead Electrothermal Digital Mel- Temp apparatus Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.11 Department of Chemistry.  327-2. Related Techniques and Instruments Technique: Removing Solvent After extraction has been completed and the extract (the solution containing the desired compound) dried, the solvent in the extract is usually removed to recover the desired compound. As a matter of fact, in an organic laboratory, there will be many situations where you want to get rid of solvents and will use one of the following evaporation techniques. A small amount of solvent can be removed by evaporation on a steam bath in a hood or by blowing it off with nitrogen or dry air in a hood. However, boiling away large amounts of solvent into the atmosphere is prohibited by environmental law. Large amounts of solvent should be removed by distillation or rotary evaporation so that the solvent can be collected and perhaps recycled. The reason the apparatus is rotated is twofold: to spread the liquid over a large surface area for more efficient evaporation, and to provide a swirling motion so that large bubbles do not form and cause the liquid to splash. Most of the apparatus consists of traps that protect the material you are saving in the round bottom flask. A bumping trap is to catch any liquid that may splash up. So that you can recover this liquid if you need it, the trap should be reasonably clean before you attach your flask. A condenser containing ice (or dry ice) condenses the evaporated solvent, which is then collected in the receiving flask for eventual discard. There may also be a trap connected between the condenser and the vacuum to protect against water backup from the aspirator, which will happen if you turn off the water before you break the vacuum – the water is sucked in. This problem can also occur, through no error of your own, if the aspirator is defective. Instrument: Rotary Evaporator The rotary evaporator (or “rotovap”) is a rather complicated piece of equipment designed to remove large amounts of volatile solvent quickly and safely. The rotovap (Figure LF.7) is hooked up to a vacuum source (either a water aspirator or a vacuum pump, and controlled by a stopcock. As you know, under a vacuum liquid boils at a lower temperature than its normal boiling point. The rotovap allows you to boil solvent away at a moderate temperature. Evaporation also causes cooling. To counteract this, you use a heating bath under the round bottom flask containing the solution. Without the bath, the solution would get so cold that ice would form on the outside of the flask, and evaporation would slow dramatically. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.12 Department of Chemistry.  327-2. Related Techniques and Instruments Figure LF.6: Rotary Evaporator (rotovap) with labeled parts. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.13 Department of Chemistry.  327-2. Related Techniques and Instruments Instrument: Büchi Rotary Evaporator The rotary evaporator (or “rotovap”) is a rather complicated piece of equipment designed to remove large amounts of volatile solvent quickly and safely. The rotovap (Figure LF.7) is hooked up to a vacuum source (either a water aspirator, 1), and controlled by a stopcock. As you know, under a vacuum liquid boils at a lower temperature than its normal boiling point. The rotovap allows you to boil solvent away at a moderate temperature. Evaporation also causes cooling. To counteract this, you use a heating bath under the round bottom flask containing the solution. Without the bath, the solution would get so cold that ice would form on the outside of the flask, and evaporation would slow dramatically. Figure LF.7: Rotary Evaporator: 1. Power switch of rotovap; 2. Knob for adjusting rotation speed; 3. Bumping trap; 4. Heating bath; 5. Temperature control display; 6. Knob for adjusting bath temperature; 7. Power switch for heating bath; 8. Stopcock for vacuum control; 9. Condenser; 10. Receiving flask; 11. Quick action jack (up/down position); 12. Clip Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.14 Department of Chemistry.  327-2. Related Techniques and Instruments Steps for Using the Rotary Evaporator 1. If a warm water bath is needed, set up the heating bath. It should be ready for use. If it is switched off: Fill the heating bath (about half-full) with water. Switch on the heating bath. The set temperature appears in the display. Do NOT change the setting, which is usually set at the optimal distilling temperature for the solvent you use. 2. Fill the condenser with ice. 3. Check the receiving flask (10 in Figure LF.7). If the flask is more than half-full with liquid from former users, ask your TA to empty it before proceeding. 4. Check the vacuum control stopcock (8 in Figure LF.7) and make sure it is open to the air (“up” position). 5. Attach your flask to the bumping trap (3 in Figure LF.7). Using plastic clip (12 in Figure LF.7), clip the flask in place with caution, keeping your right hand underneath your flask. 6. Check the main switch of the rotovap. It should be left on during the lab. In case the rotovap is switched off by the previous user, switch it on. 7. Check the water aspirator and make sure the water is turned on. 8. Gently push the down sign (▿) on the handle of the quick-action jack (11 in Figure LF.7) to lower your flask. Adjust the height of the flask until the level of solution in the flask about equal with the water level in the heating bath. 9. Start the rotation by turning the adjusting knob for rotation speed (2 in Figure LF.7). 10. Check the pump (the vacuum source) and make sure it is on. If a water aspirator is used instead, turn on water. item Turn vacuum control stopcock (8 in Figure LF.7) to vacuum source (“down” position). 11. When you have finished evaporating and are ready to remove your flask, turn the vacuum control stopcock (8 in Figure LF.7) open to air (”up” position) so that the vacuum is broken. 12. Stop the rotation by turning the rotation speed adjusting knob (2 in Figure LF.7) to 0. 13. Raise the flask gently pushing the up sign (△) on the handle of the quick-action jack 14. Unclip and remove the flask with caution. 15. You do not need to turn off the pump or the rotovap. Leave them on for the next user. If a water aspirator is used, turn off the water. Laboratory Manual Prepared by Catalyst Education, LLC for Stony Brook University LF.15 Department of Chemistry.