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CHEM 130 NAME: _____________________________ Prelab: Gas Laws and Stoichiometry Goals: 1) To utilize the Ideal Gas Law and stoichiometry to calculate the mass of a metal that reacts with an acid after experimentally determining the volume of hydrogen gas formed. 2) To learn how to use a gas buret for collection of a gas formed in a reaction. 3) To gain experience using Dalton’s Law of Partial Pressures and water vapor pressure tables. Experiment: The reaction being studied in this experiment is the metal magnesium with aqueous hydrochloric acid. HCl is a strong acid, giving the following net ionic equation: Mg (s) + 2 H3O+ (aq) --------> Mg2+ (aq) + H2 (g) + 2 H2O (l) This reaction was performed and studied in the Empirical Formula lab (#7), but this lab will focus on collection of the H2 gas formed, and with that, the mass of magnesium metal that reacted (assuming HCl is in excess) can be calculated using the Ideal Gas Law and stoichiometry. A Mg ribbon will be held in a gas buret containing the HCl and inverted in a beaker of water. The H2 gas formed displaces water in the top of the buret, allowing the volume of gas to be determined. By measuring the temperature and atmospheric pressure (using a barometer), the moles of H 2 gas produced can be calculated using the Ideal Gas equation. R is the gas constant, 0.08206 L atm/mol K. PV = nRT so n = PV / RT Dalton’s Law of Partial Pressures must be applied when determining the pressure of H 2 gas, since water vapor is also present in the gas mixture in the buret. The water vapor present can be found in a table provided, as long as the temperature is known. So the pressure of H 2 gas would be (assuming Ptotal is equal to the atmospheric pressure measured from the barometer): PH2 = Ptotal – PH20 Once the moles of H2 are calculated, stoichiometry can be applied to determine the mass of Mg metal reacted, as long as the balanced equation is known (provided above). Just use mole ratios to convert moles of H2 to moles of Mg, then use the atomic mass of Mg to convert from moles to mass in grams. There is a second method that can be used to determine the moles of H 2 gas produced. If the P and T of the gas are adjusted to Standard Temperature and Pressure (STP) conditions (i.e.: 0 oC and 1 atm), thus altering the experimental volume proportionally, then the fact that 1 mol of gas is equivalent to 22.414 L at STP can be applied. To do this, use the combined Ideal Gas Law for changing conditions to solve for the volume of H 2 gas at STP conditions (VSTP): Pexp Vexp = PSTP VSTP Texp TSTP Then just divide this volume at STP by 22.414 L/mol to get moles H 2. Stoichiometry is then applied as before to determine the mass of Mg reacted. This can be done is separate steps or in one long unit equation. CHEM 130 NAME: _____________________________ Prelab: Gas Laws and Stoichiometry Prelab Calculations to be completed prior to Lab: For the reaction: Mg (s) + 2 H3O+ (aq) --------> Mg2+ (aq) + H2 (g) + 2 H2O (l) 1) If the vapor pressure of water at 20.0 oC is 17.535 torr, and the atmospheric pressure measured by a barometer was 757.3 torr, what is the partial pressure of H 2 in the gas mixture in a buret? 2) If the volume of H2 gas collected was 47.92 ml at the temperature and pressure given in problem #1, then calculate the moles of H2 gas produced (be sure to convert P, T and V to the same units found in R!!). 3) Now calculate the mass of Mg metal reacted using stoichiometry and the balanced reaction above. 4) Determine the mass of Mg reacted using the unit equation approach involving converting the volume of H 2 gas to STP conditions, then to moles H2 and then mass of Mg. Use the experimental T, P and V given in #1 and #2. Volume H2 at STP = Mass Mg reacted =