1. While investigating the chemical and physical properties of metallic element E, a student conducts an experiment by combining E(s) and HBr(aq), which react according to the equation 2 E(s) + 2 HBr(aq) -> 2 EBr(aq) + H2(g). (a) Write the net ionic equation for the reaction. (b) In the reaction, is E(s) being oxidized or reduced? Justify your answer in terms of oxidation numbers (oxidation states). The student conducts 3 trials. The experimental apparatus is shown below. In each trial, a 10.00 mL sample of HBr(aq) is added to different quantity of E(s) in a closed test tube. To start each trial, the gas-collecting tube is filled completely with distilled water. The samples of E(s) react with HBr(aq), and the H2(g) produced is collected in the gas-collecting tube. (c) The enlarged view of the gas-collecting tube at the end of trial 3 is shown in the diagram above. What should the student record as the volume, in mL, of H2(g) collected in trial 3? (d) In the laboratory, the temperature is 295 K and the total pressure in the gas-collecting tube is 748.6 mm Hg. The vapor pressure of water at 295 K is 19.8 mmHg. Determine the pressure of the H2(g) in the gas-collecting tube.
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This is according to Dalton's law of partial pressures which states that the total pressure exerted by a mixture of gases is the sum of the pressures that each gas would exert if it were alone. So, we can write the equation as follows: Total pressure = Pressure Show more…
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This lab is based on Gas Laws, namely, the Ideal Gas law which is: PV = nRT Where P is the pressure of the gas, V is the volume of the gas, n is the number of moles of the substance, R is the Gas Constant which is 0.0824 L atm mol-1 K-1, and T is the Kelvin temperature. If you are using the above value for R, then your pressure must be in atm and volume in liter. The above equation can also be expressed as: PV = (m/M)RT Where m is the mass of the substance in grams used in the experiment and M is the molar mass of the substance. In this lab, the molar mass of the substance you are determining is magnesium. Magnesium is a solid metallic element and obviously not a gas, so how come its molar mass is determined by the gas law? The answer is simple, it produces hydrogen gas when treated with an acid as follows: Mg (s) + 2 HCl (aq) → H2 (g) + MgCl2 (aq) You can also use sulfuric or nitric acid instead of hydrochloric acid. Experiment: The experiment involves using a piece of magnesium ribbon, weighed accurately and treated with HCl. The magnesium ribbon is immersed in the acid solution in a graduated cylinder or a eudiometer. It is then inverted slowly under a trough of water. The liberated hydrogen gas is collected by the downward displacement of water. The inside pressure of the hydrogen gas is then equilibrated with the outside pressure (Boyle's law). The volume of hydrogen gas is accurately measured. The partial pressure of water vapors (Dalton's law of partial pressure) should be taken into account while calculating the pressure of the hydrogen gas collected over water because the pressure measured is actually the sum of the pressure due to the hydrogen and the partial pressure of water vapors. As shown in the balanced equation above, one mole of magnesium would produce one mole of hydrogen gas. Please read the chapter on the gas laws in your lecture textbook and watch the video. Questions: 1. When the students entered into the lab to begin their experiment, it was found that there was no magnesium in the stock room. The professor then asked the students to use any other metal instead of magnesium, however using the same procedure. Assuming the stock room has all other metals, which metal would you use in your experiment instead of magnesium? And why? 2. What are the two sources of error in this experiment that can affect your result? Explain: 3. a) What mass of magnesium would produce 53.7 L of hydrogen at 22°C and 754 mm Hg to give a molar mass of 23.6 g/mol (determined experimentally)? b) If the accepted molar mass of magnesium is 24.31 g/mol, calculate the percent error above:
Sri K.
2. A sample of oxygen at 22.0 °C and 740 mmHg was found to have a volume of 435 mL. How many grams of O2 were in the sample? 3. A sample of argon is trapped in a gas bulb at a pressure of 760. mmHg when the volume is 150.0 mL and the temperature is 38.0 °C. What must be its temperature if its pressure becomes 715 mmHg and its volume is 212.0 mL? 4. A student has prepared some CO2 by heating CaCO3(s) to high temperatures: CaCO3(s) → CO2(g) + CaO(s) a. If a volume of 625 mL of CO2 was produced at a pressure of 735 torrs and a final temperature of 25.0 °C, find the mass of CO2? 5. Ethylene gas burns in air according to the following equation: C2H4(g) + 3 O2(g) → 2 CO2(g) + 2 H2O(l) If 75 L of C2H4 at 25°C and 1.085 atm burns completely in oxygen, calculate the mass of CO2 produced. 6. A 1655 mL container of oxygen gas is at 28 °C and 748 torrs. Hydrogen gas is pumped into the container, producing water. What is the least amount of mass of hydrogen gas needed in order to react the oxygen to completion? 2 H2(g) + O2(g) → 2 H2O(l) 7. What is the volume of oxygen at STP when 200 g of potassium chlorate decomposes to form potassium chloride and oxygen? 2 KClO3 → 2 KCl + 3 O2 8. What is the volume of carbon dioxide produced by the combustion of 45 g of ethanol C2H5OH at 55°C and 725 mmHg? C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(l)
Explain why it is necessary to bring the water level inside the graduated cylinder equal to the water level outside the graduated cylinder after gas collection is complete. Which variables of the Ideal Gas Law are likely to be affected if you don't (and how)? If calcium metal is used as an unknown in an experiment of this type (and it may be), what would be a reasonable value for 'n', assuming the sample mass was 0.063 g? Write the balanced chemical equation for the reaction of calcium metal with aqueous hydrochloric acid. Using the value determined above, determine the amount of hydrogen gas produced in mL. Assume the temperature and pressure to be the same as they were on the day you did the experiment in the lab. Explain how failing to account for any trapped air inside the graduated cylinder at the start of the experiment affects the calculation of: (a) Volume of H2 gas produced, (b) Number of moles of H2 gas, (c) Molar mass of the unknown metal sample. What changes, if any, would you make to the experimental procedure if you decided to use aqueous sulfuric acid (H2SO4) instead of aqueous hydrochloric acid (HCl)? Write the balanced chemical equation for the reaction of the unknown metal you were given with aqueous sulfuric acid.
Adi S.
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