Determining the Molar Mass of a Volatile Liquid

Determining the Molar Mass of a Volatile Liquid Introduction In the world of chemistry, it is crucial to have the ability to determine specific properties of unknown substances in order to classify and categorize them. These properties include, but are not limited to: the molar mass of a substance, the molar volume, density (of a gas, liquid, or solid), partial pressure of a gas, and boiling points. In this experiment, the main focus was to determine the molar mass of a volatile liquid after vaporization and re-condensation has been performed. When a gas behaves ideally, it acts in a certain way that, although does not comply with real gaseous properties, accurately gives information of the quantitative properties of the gas. These are related through the ideal gas law: PV = nRT Where P is the pressure, V is the volume of the gas, n is the number of moles, R is a gas constant, and T is the temperature of the gas converted to Kelvin. It is a well-known fact that the density of a gas is significantly lower than that of a solid or a liquid, simply because gaseous molecules do not stick together and completely fill the volume of the container that they are in. Therefore, there is less mass per unit of volume of the substance in the container. The Dumas method involves condensing a container filled with gas into a liquid so that the mass of the liquid can be weighed and correlated with the number of moles of gas that were produced within the flask. Since gases fill the volume of the container in which they are placed, an effusion hole must be added to the container so that the gas may completely fill the volume and allow excess vapor to escape. This allows for more accurate results when determining the specific properties of an unknown substance. Procedure Despite being a simple experiment, there are minute details that, if missed, will alter the finals results. To begin, a 250 mL Erlenmeyer flask was properly cleaned, dried, and weighed on a scale where its exact mass was weighed to the nearest 0.01 gram. It was crucial to ensure that all water molecules were removed from the flask to promote accuracy. A 1000 mL beaker was placed on a hot plate and filled with water to begin heating, but not to the boiling point yet. The dry Erlenmeyer flask was clamped to a ring stand and submerged in the water until the beaker was within a centimeter of the rim. Two boiling chips were dropped into the heating water to promote even boiling and approximately 5 mL of unknown liquid # 802 were measured into the flask. The hot plate was then turned up to increase the temperature of the water to begin the boiling process. A piece of aluminum foil was fitted tightly over the cap of the Erlenmeyer flask where a tiny pinhole was made in the center to allow excess vapor to effuse. It was crucial to ensure that no vapor could escape from the flask at any point except for the hole. While waiting for the unknown liquid to completely evaporate and reach the same temperature as the boiling water bath (100 degrees Celsius), a pneumatic trough was filled a little over halfway with cold water which was subsequently used to condense the vapor produced by the boiling water. Three minutes after the liquid had completely evaporated, the inside of the Erlenmeyer flask was at equilibrium with the temperature of the boiling water and the flask was quickly transferred to the cold water bath using paper towels to prevent accidental burns. After about 8-10 minutes, all of the water was thoroughly dried from the flask, which was then taken to the weighing station where the aluminum foil was removed and the flask containing the condensed liquid was weighed and recorded to the nearest 0.01 gram. In order to find the exact mass of the vapor that filled the flask, the initial mass of the dry flask was subtracted from the final mass of the flask containing the condensed liquid. The resulting value was equivalent to the mass of the vapor that filled the Erlenmeyer flask. After gathering these results, the experiment was repeated using 3 mL of the unknown liquid. The resulting mass, which was relatively the same as the first mass found, was averaged out to find the average mass of vapor at 100 degrees Celsius and approximately 1.00 atm. All of the condensed liquids were discarded in the proper lab waste container after all of the data was collected. To find the total volume that the 250 mL Erlenmeyer flask can contain, the flask was completely filled with tap water, which was transferred to a graduated cylinder so that the total volume could be recorded. Using a local barometer, the atmospheric pressure for the day was established and used in the ideal gas equations. Results Note: Bolded values are important for calculations Table 1: Weight of condensed liquid in the Erlenmeyer flask Unknown Liquid Dry, empty 250- Trial 1 5 mL 118.35 grams Trial 2 3 mL 118.35 grams Average 4 mL 118.35 grams mL flask Flask containing 118.46 grams 118.48 grams 118.47 grams condensed liquid Mass of 0.11 grams 0.13 grams 0.12 grams Unknown Liquid Eq. 1 – Finding Mass of Unknown Liquid Final Mass – Initial Mass = Mass of Substance Table 2: Standard information regarding experiment and surroundings Temperature of Water Temperature of Vapor (eq. 2) Atmospheric Pressure (eq. 3) Volume of Erlenmeyer Flask 100 degrees Celsius, 373.15 Kelvin 100 degrees Celsius, 373.15 Kelvin 29.74 inHg, 0.9939 atm 0.2796 L Eq. 2 – Converting degrees Celsius to Kelvin T (Kelvin) = T (Celsius) + 273.15 Eq. 3 – Pressure conversions 1 inHg = 2.54 cmHg 1 cmHg = 10 mmHg 1 atm = 760 mmHg Table 3: Information about unknown liquid # 802 Molar Volume (eq. 4) Density of Vapor (eq. 5) Molar Mass of Substance (eq. 6) Moles of Substance (eq. 7) 30.81 L/mol 0.43 g/L 13 g/mol 0.009075 mol Eq. 4 – Finding the molar volume of a substance V = (RT)/P V = molar volume in L/mol, R = gas constant, T = temperature in Kelvin, P = pressure in atm Eq. 5 – Finding the density of a gas two different ways D = m/v = (PM)/(RT) D = density, m = mass, v = volume of container that the gas fills, M = molar mass Eq. 6 – Finding the molar mass of a substance by altering the above equation M = (mRT)/(PV) Eq. 7 – Finding the number of moles of the substance in the container PV = nRT Plug in all known values during the experiment and solve for n to find the number of moles. As it turns out, the unknown liquid used was methanol, which has a molar mass of 32.04 g/mol. The next section will discuss what may have resulted in such a large experimental error. Eq. 8 – Finding percent error: % error = |32.04 – 13|/32.04 x 100% = 59% error Discussion Several things can go wrong in any experiment and ultimately result in inaccurate data. It is crucial to minimize these experiment and systematic errors by first analyzing what could go wrong. When initially weighing the flask, ensure that it is completely dry and does not contain any excess water molecules that may affect the mass recorded. Weighing the flask several times will ensure that an accurate reading is being given. When a pinhole it poked through the aluminum foil that covers the mouth of the flask, the hole must be made as tiny as possible to reduce an increased rate of effusion. The sole purpose of this hole is to allow any extra vapor to escape the flask until the volume of the flask is completely filled with the gas. Whenever there is no remaining liquid in the bottom of the flask, it is assumed that the volume of the vapor is exactly equal to the volume of the flask, which is reported in table 2. As the heating process occurs, molecules in the volatile liquid gain kinetic energy and eventually (at their boiling points) overcome the potential attractive forces keeping them in liquid form and escape into the gas phase. Approximately three minutes after all of the liquid has escaped into the gas phase, the inside of the Erlenmeyer flask will be isothermal with the boiling water on the outside, at 100 degrees Celsius. This is the temperature measurement that is to be used in the equations listed in the data section. Whenever the flask, completely filled with gas particles, is transferred over to the ice cold water bath, the gas molecules begin to lose kinetic energy and eventually convert back to the liquid form. In short, the same process occurred in the opposite direction. The weight of the newly formed condensed liquid inside of the flask can be calculated by subtracting the initial mass from the final mass (eq. 1). This experiment should be repeated and end with matching results to ensure that it was carried out correctly. The molar mass that was calculated after the experiment stands at a 59% error. It is difficult to pinpoint the exact reason for this mishap, but several theories can be made. It is assumed that the gases in this experiment behave ideally, meaning that they do not interact with one another and they all behave the same. However, this is untrue to how real gases behave and can cause deviations. These ideal gases can be measured most accurately at low pressures and high temperatures, such as those that occur in this experiment. There may be more efficient ways to identify an unknown substance, such as determining its mass in a certain volume and finding the density. This can be compared with a list of densities of substances in which the correct one can be identified. If the substance were in solid form, it could be titrated with a corresponding acid or base until all of the substance has been consumed by a chemical reaction (often shown via an indicator) and the moles of substance could be calculated using the chemical equation. Another method of identifying a substance is to test its physical properties such as boiling point, melting point, viscosity, and surface tension. Molecules with more powerful intermolecular forces will exhibit higher levels of these physical properties because it is harder to separate the individual atoms. This experiment can easily be repeated with any volatile liquid to determine its quantitative properties, such as molar mass, so that it may be identified. It does not matter how much of the liquid initially goes into the flask because it will all be evaporated out and any of the excess vapor that cannot remain in the volume of the flask can effuse through the pinhole. If this method is used to find the properties of a liquid that has a higher boiling point than water (100 degrees Celsius), then a different liquid will have to be used to heat up the flask. As seen in table 3, the density of a gas is extremely small (which is why it was reported in g/L rather than g/mL) and this density corresponds with the amount of vapor that completely filled the volume of the container. Conclusion Through studying these concepts, it is clear that there are numerous ways to determine a substance through means of finding variables under certain conditions and importing them into known equations. Although a gas is significantly less dense than its liquid form, the moles of vapor that completely filled the volume was exactly the same as the number of moles of the resulting liquid. This shows that although a phase change occurred, the amount of the substance remained constant in its transition from a gas to a liquid. Experiments like this could be used to extensively study the changes that occur during phase transitions and also the similarities between two different phases of the same substance. By carefully examining the substance on a molecular level during a phase change, one can determine the amount of energy needed for a molecule to overcome its attractive forces, and what must occur for it to revert back to its original form. These studies help the scientific world better understand the relationship between the different phases of a volatile substance.