2 lab periods; work in pairs. Complete the Preparation page before laboratory.
Goals
Background
It is not an exaggeration to say that all chemical and physical interactions are a result of the attraction of positive charges for negative ones. The fundamental forces among such particles account for the existence of the atom; for the bonding of atoms to form molecules; for the shapes adopted by these molecules; and for the interaction/reaction between one molecule and another. When two molecules come together, the positive portion of one molecule seeks the negative portion of the other. The consequence of this is that chemical reactions take distinct and to some extent predictable paths.
From very simple information about the percent composition of a covalent compound, usually determined at least in part by combustion analysis, it is possible to determine its empirical formula. When this is combined with molecular mass information, obtainable by a variety of means, the molecular formula, showing the actual numbers of atoms of each type present in a molecule of the substance, may be deduced. Reasoning then from the preferred bond numbers for each of the atoms, which follow from the electron configurations of the atoms, it is possible to develop a Lewis structure for the molecule. A Lewis structure is a picture that attempts to represent the bonding structure and electron distribution in the molecule. Once the Lewis structure is at hand, the principles of VSEPR theory may be applied to provide a description of the 3-dimensional shape of the molecule. From the 3-D structure and a table of atom electronegativities, bond polarities may be assigned and combined to determine whether or not the molecule has a dipole moment; i.e., whether the molecule has positive and negative ends. Once we have an understanding of the electrical nature of the molecule, we are in a position to predict (in some cases) or to rationalize (in most cases) the properties, both physical and chemical, of the substance. Clearly, then, the ability to arrive at an understanding of the polarity (or lack thereof) is crucial, because polarity gives rise to the forces that dictate the properties and behavior.
ALL forces at the atomic/molecular scale are electrical in origin, and can be fundamentally understood in terms of Coulomb's Law:
(1): F = -K(q1)*(q2)/r2
Here q1 and q2 are the magnitudes of two interacting charges, r is the distance between them, and K is a proportionality constant. If both charges are positive or both are negative, a negative force results, which we interpret to mean that the particles repel each other and tend to move apart. If one charge is positive and the other negative, a positive force results, which we interpret to mean that the particles attract one another, and tend to move together.
The major types of forces that operate among atoms, molecules, and ions are indicated in Table 1.
| Type | Example |
|---|---|
| Covalent chemical bonds | C-Cl bond in CCl4 |
| Ionic chemical bonds | Bond between Na+ and Cl- in NaCl |
| Metallic chemical bonds | Bond between Na atoms in Na(s) |
| dipole-dipole | Force between 2 like or unlike polar molecules such as two acetone molecules or an acetone molecule and a chloroform molecule |
| Hydrogen bonds (special type of dipole-dipole) | Force between 2 water molecules |
| ion-dipole | Force between a Na+ ion and a water molecule |
| dipole-induced dipole | Force between a water molecule and a CO2 molecule |
| instantaneous dipole-induced dipole (dispersion) | Force between two nonpolar molecules of I2 |
This rather small collection of types of forces forms the basis for all of chemistry. We can briefly discuss the types of force, roughly in order of decreasing strength.
Chemical Bonds. Chemical bonds can be ionic (transfer of electrons), covalent (sharing of a pair of electrons between two atoms), or metallic (sharing of electrons among huge numbers of atoms). Covalent bonding often leads to the formation of small (or large) aggregates of atoms called molecules, which function as a single unit. Ionic bonding leads to the formation of essentially infinite arrays of positive and negative ions, held together by the strong attractive Coulomb forces exerted by oppositely charged ions on one another. In an ionic compound such as NaCl it is not possible to distinguish a unit of NaCl, since each Na+ ion is surrounded by 6 Cl- ions, each of which is in turn surrounded by 6 Na+ ions, and on and on. Small molecules can not be identified in ionic compounds. For this reason we consider the "unit" in an ionic compound to be the entire crystal, consisting of huge numbers of + and - ions. Metallic bonding will not be discussed here.
Dipole-dipole Forces. If a chemical bond involves two different atoms, or even two atoms that are the same but are in different environments in the molecule, then the bond is polar, because the pair of bonding electrons is drawn closer to the more electronegative atom of the bond. We represent the bond polarity by drawing a vector along the bond pointing from the more positive to the more negative atom. A molecule containing several polar bonds may or may not be polar overall, depending upon how the polar bonds are oriented in space. If the polar bonds are oriented so as to cancel (negate) each other, then the molecule will have no NET polarity. On the other hand, if the polar bonds are oriented so as to reinforce each other, then the molecule will have net polarity. Such a molecule is said to have a molecular dipole moment. The molecular dipole moment is the vector sum of the bond dipoles. Of course, it is also a vector. Two proximate polar molecules tend to orient so that the negative end of one molecular dipole is close to the positive end of the other. The resultant electrical force of attraction is called a dipole-dipole force. Dipole-dipole forces operate between neighboring molecules of acetone in liquid acetone, and similarly in many other substances.
Hydrogen Bonds. Molecules containing one or more hydrogen atoms attached directly to either nitrogen, oxygen, or fluorine exhibit a particularly strong dipole-dipole interaction called hydrogen bonding. The bond involving this hydrogen atom is strongly polar, resulting in a substantial positive charge on the hydrogen atom. The positive end of this dipole is strongly attracted to the negative end of a similar dipole on an identical or similar molecule. Because the hydrogen atom is so small, the attracting ends of the 2 dipoles can approach closely, leading to a strong force. Hydrogen bonding is responsible for many of the properties of liquid water, including but not limited to its unusually high boiling point, its effectiveness as a solvent, and the relative densities of its solid and liquid phases. Hydrogen bonding is profoundly important in biochemistry.
Ion-dipole forces. Many ionic compounds dissolve readily in water and in other highly polar solvents because the positive and negative ions are attracted respectively to the negative and positive ends of the dipole of the water molecule. The resulting ion-dipole forces allow the separated + and - ions to be stable in aqueous solution. Clearly, ion dipole forces can be important only in mixtures involving 2 substances, one ionic and one polar covalent.
Dipole-Induced dipole forces. If a covalent substance with polar molecules is mixed with a covalent substance with non-polar molecules, one might expect little attractive force between them. However, it is possible for the dipole of the polar molecule to influence the electron distribution in a nearby non-polar molecule so as to induce a temporary dipole in the nonpolar molecule. The polar and nonpolar molecules then attract each other, and so have some tendency to stick together. This type of force is operative when a polar substance such as acetone is mixed with a non-polar substance such as carbon tetrachloride.
Instantaneous dipole-induced dipole (dispersion) forces. Even nonpolar molecules are capable of exerting electrical forces on one another. We know this because many substances with nonpolar molecules (e.g., I2, CCl4, C6H6) are liquids or even solids under ordinary conditions. They could not exist in these phases unless some attractive forces were operative. A way to understand the nature of these forces is to remember that the electrons in molecules are mobile--i.e., they are in constant motion. At any particular instant, the electrons in a nonpolar molecule can move (accidentally) somewhat in concert toward the same side of the molecule. Because the nuclei do not move, this creates what is called an instantaneous dipole in the molecule. This is shortlived, but persists long enough to induce weak dipoles in neighboring molecules, which are then mutually attracted. The resulting force is called the instantaneous dipole-induced dipole force, or dispersion force for short. It is dispersion forces that are responsible for the ability of nonpolar molecules to visit the liquid and solid states. Because all molecules, polar or nonpolar, have electrons, we expect that dispersion forces will be operative among all molecules, not just non-polar ones. They are just most obvious in the nonpolar ones, in which no other forces operate.
Collectively, dipole-dipole, dipole-induced dipole, and dispersion forces are sometimes called intermolecular forces because they result from the attraction of one covalent molecule for another. If we agree to use the phrase, interunit forces rather than intermolecular forces, we can include ion-dipole forces as well. In this project, you will explore interunit forces in a number of ways. First, you will prepare mixtures of a number of pure substances in order to explore molecular forces. You will determine whether or not the formation of mixtures leads to temperature and volume changes in the system. From these very simple measurements of DT and DV, you can follow a clear line of reasoning back to the forces operating between molecules before and after mixing. For clarity in development, we base the thinking on a system in which T increases and V decreases when the mixture is formed.
Chain of Reasoning:
Second, you will learn and use a technique for measuring the boiling point of a small volume of liquid. Boiling point is an indication of the strengths of interunit forces in the liquid phase.
Third, you will measure vapor pressures of several liquids at several temperatures to determine the depth of the liquid potential well. The variation of vapor pressure with temperature is given by the Clausius-Clapeyron equation, (2), where Pvap is in atm, T is in K, DHovap the standard enthalpy of vaporization in J/mole, DSovap the standard entropy of vaporization in J/K-mole, and R the gas constant.
(2): ln Pvap = -DHovap/RT + DSovap/R
A plot of ln Pvap vs. l/T should be linear with slope related to DHovap and intercept to DSovap.
(3): Slope = -DHovap/R
(4): Intercept = DSovap/R
Equation (2) can be used to calculate the normal boiling point of the liquid.
Finally, you will explore the interesting phenomena of surface tension and capillary action.
Focus Questions
Required Materials
Safety
Wear safety goggles at all times in the laboratory. Most organic substances (e.g., ethanol, isopropanol, acetone) are flammable; keep away from open flames. Concentrated acids and concentrated solutions of bases burn the skin. Should a spill occur, flush liberally with tap water. Chromium is a known carcinogen. Avoid contact of Na2Cr2O7 with the skin.
Experimental
Carry out the following procedures. Carefully record all data, observations, results, and conclusions in a laboratory notebook.
Liquid-Liquid Mixtures. Carefully add liquid A (see Table 1) to a 10-mL graduated cylinder until the liquid meniscus is at exactly the 5.00 mL mark. Measure and record its temperature, and withdraw the thermometer. Check to make sure that the volume level in the cylinder does not change significantly when the thermometer is withdrawn. Using a 5-mL graduated pipet, draw up 5.00 mL of liquid B. Assuming that the bottles containing liquids A and B have been sitting near each other in the lab, they should be at the same temperature. Eject liquid B from the pipet to the graduate containing liquid A. This must be done forcefully enough to result in mixing of the liquids, but not so forcefully that splashing occurs. Immediately after mixing, record the total volume of the mixture. Then stir the mixture with the thermometer and measure the final temperature of the mixture. Record any other observations that you consider pertinent. After doing measurements on a particular pair, dispose of the liquid mixture appropriately, thoroughly dry the graduate, and proceed to the next pair. Do the water/dichloromethane mixture last, and save it. IN DOING THESE EXPERIMENTS, YOUR FOCUS SHOULD BE ON DT (THE CHANGE IN TEMPERATURE THAT OCCURS ON MIXING) AND ON DV (THE CHANGE IN VOLUME THAT OCCURS ON MIXING).
| Liquid A | Liquid B | Dispose of mixture in |
|---|---|---|
| H2O | H2O | sink |
| Isopropanol | non-halogenated organics | |
| Acetone | non-halogenated organics | |
| Dichloromethane | SAVE | |
| 6 M H2SO4 (add this SLOWLY TO
water, not the reverse. DO NOT touch the graduate after mixing) | sink, with water flush | |
| Concentrated NH3 (add this TO water, not the reverse) | sink, with water flush | |
| Ethanol | acetone | non-halogenated organics |
| Dichloromethane | halogenated organics | |
| Concentrated NH3 (add this TO ethanol, not the reverse) | non-halogenated organics | |
| Acetone | dichloromethane | halogenated organics |
| heptane | non-halogenated organics |
Into a 1-dram vial, pipet 0.5 mL of distilled water and 0.5 mL of dichloromethane. Cap the vial and shake. Make careful observations. Then add ethanol to the vial dropwise, 20 drops at a time. After addition of each 20 drops, cap the vial and shake. Make careful observations. Add a total of 100 drops of ethanol. Dispose of the mixture in the halogenated organics waste.
Solid-Liquid Mixtures. Weigh 1.0 g of each solid, B, below. Using a graduated pipet, measure 5.00 mL of liquid A and transfer it to a small beaker. Measure its temperature. Add solid B to liquid A, with stirring. Measure the temperature of the mixture. Make careful observations and record them. Again your focus should be on the DT resulting from mixing.
| Liquid A | Solid B | Dispose of solution in |
|---|---|---|
| H2O | NaCl | sink |
| NH4NO3 | sink | |
| NaOH | sink | |
| sugar | sink | |
| Acetone | NaCl | non-halogenated organics |
| NaOH | non-halogenated organics | |
| sugar | non-halogenated organics | |
| 2 small crystals I2 | non-halogenated organics | |
| Dichloromethane | NaCl | halogenated organics |
| sugar | halogenated organics | |
| 2 small crystals I2 | halogenated organics |
Divide the saved water-dichloromethane mixture in half (how do you propose to do this?) To one part, add a few small crystals of I2(s), with stirring. Make careful observations. To the second part, add a few crystals of Na2Cr2O7, with stirring. Make careful observations.
Discard both mixtures in halogenated organic waste.
Vapor Pressure by Rate of Evaporation. Obtain a glass microscope slide and an automatic pipettor equipped with a plastic tip. In turn, draw up 5 microliters of each of the following liquids, eject the liquid onto the slide, and measure the time required for complete evaporation of the liquid. Do at least 3 runs for each liquid.
When finished, discard any unused organic solvents in the appropriate waste bottles.
Vapor Pressure by a Syringe Method. Obtain a small side-arm test tube, a 10-mL glass syringe, a septum cap to fit the mouth of the test tube, and a 1-mL plastic syringe. The sidearm test tube and syringe must be absolutely CLEAN and DRY. Accomplish this for the syringe by rinsing the inside wall of the barrel and the surface of the plunger thoroughly with acetone. Wipe the entire surface of the plunger and dry the inner wall of the barrel with a Kimwipe. Insert the plunger and move it in and out several times. Withdraw it and let acetone evaporate from its surface. Repeat the insertion and withdrawal until the acetone is gone. At this point, the plunger should slide and rotate freely, with almost no friction. Get it to this point before proceeding. Once the equipment is clean and dry, remove the glass syringe plunger and lay it on the bench top while you set up the apparatus in Figure l. Attach a short length of appropriately sized rubber tubing to the end of the glass syringe. The fit must be tight. Attach the other end of the tubing to the sidearm of the test tube. This fit must also be tight. Use ring stands and clamps to mount the test tube vertically and the syringe horizontally. BE SURE THAT THE SYRINGE IS EXACTLY HORIZONTAL SO THAT WHEN THE PLUNGER IS REPLACED, IT WILL NOT FALL OUT AND BREAK. DO NOT TIGHTEN THE CLAMP ON THE SYRINGE BARREL--INSTEAD, ALLOW THE SYRINGE TO REST ON THE LOWER ARM OF THE CLAMP. Then replace the plunger in the barrel of the syringe. Fit the solid end of the septum cap tightly into the top of the test tube and fold the hollow end over the lip of the tube. Insertion of the septum cap will probably affect the position of the syringe plunger. Write down the plunger position.
In the hood, pour 1-2 mL of CH3OH (methanol) into a 1-dram vial, and cap the vial. Take it to your work area and place it on a labelled paper towel strip.
Now use the small plastic syringe to inject about 0.1 mL of CH3OH into the sidearm test tube. To do this, you will need to puncture the center of the septum cap with the syringe needle. Once the liquid is in, remove the small syringe from the septum cap, and observe the position of the plunger of the glass syringe. Write down the final plunger position. Remove the septum cap from the test tube, and aspirate the test tube to remove ALL TRACES of CH3OH, both liquid and vapor. While aspirating, move the syringe plunger in and out at least 10 times to expel all methanol vapor from the syringe.
Repeat the measurement of the vapor pressure of pure liquid until you are confident of the reproducibility of the value to within 10%.
Transfer 1.00 mL of CH3OH to another 1-dram vial. Add 0.10 g benzophenone, cap, and swirl or shake to dissolve the solute. This may take a few minutes. Then measure the vapor pressure of the solution using exactly the same procedure as used above. Clean the sidearm test tube by rinsing several times with small portions of acetone, then aspirate dry. Repeat rinsing and drying until you see NO benzophenone on the walls of the tube. Repeat the experiment to obtain a reliable value for the vapor pressure of the solution. Transfer the contents of the vials to the nonhalogenated organic waste bottle. Rinse vials with acetone, and pour this in a nonhalogenated organic waste bottle.
Measure the volume capacity of the side arm test tube by inserting the septum cap, and filling the tube through the sidearm with distilled water. When the tube is completely full, pour the contents into a graduated cylinder. The volume of water in the cylinder provides a good estimate of the volume capacity of the sidearm test-tube portion of the vapor pressure assembly.
Use your data to calculate the vapor pressures of pure CH3OH and the solution of benzophenone in CH3OH, in torr, using the following equation:
Be sure to thoroughly clean and dry the sidearm test tube and syringe (if necessary).
Discard unused methanol and benzophenone solution in non-halogenated organic waste.
Temperature Dependence of Vapor Pressure.
You will use the same apparatus here as in the previous part of the project. Clean and dry the sidearm test tube and the syringe thoroughly as above. Set up an iron ring and wire gauze on a ring stand. Fill a l50-mL beaker half full of ice and add some water to make a slush. Set up the apparatus in Figure l; make sure that the syringe plunger is all the way into the barrel. Submerge the bottom 1 inch or so of the sidearm test tube in the ice bath. Attach a Bunsen burner to a gas jet and place the unlit burner under the beaker.
Using a 1-mL plastic syringe, draw up about 0.1-0.2 mL of your first assigned liquid. Holding the syringe vertically with the needle up, push the plunger into the syringe to expel all air. Catch liquid that exits the needle using a tissue. Insert the needle of the plastic syringe through the septum cap until you see it enter the upper part of the sidearm test tube. Inject the liquid, and retract the needle. Remove the plunger from the barrel of the plastic syringe and set the pieces on a Kimwipe to dry. Stir the water in the ice bath with a thermometer and measure the ice bath temperature. Record the temperature and the position of the plunger of the glass syringe.
Fire up the Bunsen burner and heat the ice bath until the ice melts. Stop heating when the temperature of the bath is 3 or 4 degrees, and stir until the temperature stabilizes. The syringe plunger may tend to stick in the barrel, so frequently rotate the syringe plunger without moving it in or out of the barrel to keep it moving freely. Record the stable temperature and the position of the syringe plunger. Briefly heat the bath again to obtain a stable temperature about 10 degrees higher than the previous one. Record temperature and plunger position. Continue at 10-degree intervals to a maximum temperature of 50 oC. When finished, turn off the burner and disassemble the sidearm test tube-syringe apparatus. Aspirate all traces of liquid and vapor from the interior of the test tube and syringe.
Repeat the entire procedure for your 2 other assigned liquids.
After disassembling the syringe-test tube apparatus, fill the sidearm test tube with water. Insert the septum cap and allow all air bubbles to escape through the sidearm. Then fill the sidearm with water. Empty the test tube into a 50-mL graduated cylinder and record the water level. This gives you a pretty good measure of the volume capacity of the sidearm test tube.
Dispose of all leftover liquids in the organic waste bottles in the fume hoods.
Boiling Points of Pure Liquids. Fill a 150-mL beaker about 3/4 full with water and support it on a ring stand using an iron ring/wire gauze. Obtain a set of 3 test tubes (1x8 in, 6x3/4 in, 6x1/2 in) and nest them. Using a triangular file, scratch a melting point capillary about 2-3 cm from one end, and break off the 2-3 cm segment. Seal one end of the capillary segment in the burner flame, and place the segment sealed-end-up into the inner tube of the nested set. Place a filter adapter on the outer nested tube, wide-end-up, leaving about 1-inch of tube above the adapter. Clamp the nested tubes above the filter adapter so that it extends about 2 cm into the water bath. Using a Pasteur pipet, add your first liquid to the inner glass test tube to a depth of about 1 cm. Then insert a digital thermometer into the inner tube. Fill the filter adapter with small pieces of dry ice. This will serve as a condenser for the solvent vapors, preventing them from exiting the test tube. YOU WILL HAVE TO KEEP AN EYE ON THE DRY ICE, AND REPLENISH IT FREQUENTLY, BECAUSE IT WILL SUBLIME RAPIDLY OVER THE HOT WATER BATH. Begin heating your water bath at a rate of 2-3 oC per minute, keeping an eye on the bottom end of the capillary segment. As the water bath warms the liquid, you will see bubbles exit from the open end of the capillary. These will be infrequent as long as the temperature is quite far below the boiling point of the liquid. (CHECK YOUR DRY ICE.) As the temperature approaches the boiling point, however, bubbling will become more rapid. At the boiling point, a steady stream of vapor bubbles will be seen exiting the bottom of the capillary. (CHECK YOUR DRY ICE.) When you see the steady stream, decrease the rate of heating of the water bath. The temperature should level off and remain constant as the liquid refluxes in the inner tube. Record the steady thermometer reading, which is the boiling point.
Remove the nested test tubes from the hot bath and allow the liquid to cool below its boiling point. (CHECK YOUR DRY ICE.) Put a second capillary segment in the inner tube, clamp the nested tubes in the bath, and SLOWLY heat the bath. You should look for a stream of bubbles from the capillary, and the thermometer reading should remain constant. The temperature should be within 0.1-0.2 degree of your previous value. Remove the nested tubes from the hot bath.
Repeat the process a third time if your first 2 boiling point temperatures differ by more than 0.2 degree. (KEEP YOUR CONDENSER STOCKED WITH DRY ICE.)
Repeat the entire procedure for your other two liquids.
Boiling Points of Solutions. Choose one of your three liquids for this part of the project. Using this liquid as solvent and an appropriate solid as solute, prepare a SMALL VOLUME of solution in which the mole fraction of solute is about 0.1. For example, you might measure out exactly 2 mL of solvent, then add the appropriate (weighed!) amount of solute. Measure the boiling point of this solution using the procedure above. When finished, discard liquid and solution in proper waste containers.
Surface Tension. The instructor will provide you with a number of items, most of them flat, to use in the surface tension experiments. Make sure all the items are completely dry, then weigh them and record their masses. Fill a clean 100- or 150-mL beaker to near the top with distilled water. Using clean tongs, carefully place in turn each of the provided items on the surface of the water. Your goal is to determine whether or not each object "floats" on water, so it is important to maintain the flat surface of the object parallel to the water surface when placing the object. In each case, carefully observe and record whether each object floats or sinks. Remove each item from the water before proceeding to the next one.
You should investigate a large enough collection of objects to determine what properties of an object determine whether or not it floats on water. Possibly relevant properties include density of material, mass of object, surface area of contact between the object and the water surface, and shape of the object.
When finished, place all solid items on a piece of paper towel to dry. Discard the distilled water.
Meniscus Formation and Capillary Action. Fill a clean 4-inch or 6-inch test tube halfway with distilled water. Insert a length of glass tubing having an internal diameter (id) of 4mm. Hold the tubing so that it rests against the bottom of the test tube and is centered in the tube. Make and record observations. Remove the length of glass tubing and insert a length of plastic tubing into the test tube. Make and record observations. Dry the glass and plastic tubing using aspiration.
Fill a clean 4- or 6-inch test tube halfway with dichloromethane. In turn, insert the length of glass tubing and the length of plastic tubing as above. Make and record observations.
Predict what would happen if you inserted the glass and plastic tubing into test tubes containing ethanol; acetone; hexane.
When finished, discard dichloromethane in the halogenated organic waste bottle. Discard the distilled water. Return glass and plastic tubing to the instructor.
Determination of Polarity. Set up 3 burets. Into the first buret put about 5 mL of water; into the second, 5 mL of acetone; into the third, 5 mL of carbon tetrachloride.
Remove all static electricity from a glass stirring rod by submerging it about halfway in a beaker of distilled water for a couple of seconds. Withdraw the rod from the water, and SHAKE excess water from the rod. All static electricity should now have been removed from the rod. It is important not to rub the rod with paper towel or Kimwipe to dry it, because this may create static electricity on the rod. In turn, open each buret stopcock so that a stream of liquid flows from it, catching the stream in a beaker. Move the glass stirring rod toward the stream of liquid and observe the effect. Record observations.
Put a nitrile lab glove on your left hand and pull the stirring rod through your gloved hand several times. This will create a static electic charge on the stirring rod. Again open each buret stopcock in turn so that a stream of liquid flows from it into a beaker. Move the charged glass stirring rod toward the stream of liquid and observe the effect. Record observations.
When finished, discard carbon tetrachloride in the halogenated organic waste bottle, and acetone in the non-halogenated organic waste bottle. Discharge the stirring rod by submersion in distilled water.
Disposal
Instructions for proper disposal accompany the separate experiments above.
