The Gram-Molecular Volume of Gases

The objective of this lab is to investigate gas laws and various calculations related to gases. Students will use tools and equipment for scientific analysis, demonstrate safe practices in the chemical laboratory, apply their understanding of gas laws, chemical reactions, molecules, compounds, and chemical composition, and use exponential notation and significant figures appropriately. They will also record their results, use scientific reasoning to evaluate physical and natural phenomena, and identify unifying themes in the scientific field of study.

The discussion provides an overview of the concepts and principles involved in the experiment. It starts by stating Dalton's law, which states that equal volumes of gases at the same temperature and pressure contain the same number of molecules. This leads to the concept of the molar gas volume, where 22.4 liters of any gas at standard temperature and pressure (STP) contains the same number of molecules as 32 grams of oxygen (O2) at STP.

The experiment aims to measure the density of oxygen by preparing a known weight of oxygen and measuring the volume it occupies. Oxygen is prepared by heating potassium chlorate (KClO3) with the presence of a catalyst, ferric oxide (Fe2O3). The loss in weight of the test tube containing KClO3 represents the weight of oxygen produced.

The experiment utilizes Dalton's law to find the pressure of the oxygen gas generated. Boyle's and Charles' laws are then applied to convert the volume of oxygen gas generated to the corresponding volume at standard temperature and pressure (STP). The underlying basis for the experiment is Avogadro's law, which states that equal volumes of gases, at the same temperature and pressure, contain the same number of molecules.

The lab procedure involves weighing the test tube and then adding KClO3 and Fe2O3 to it. The experimental apparatus is assembled, including a glass bottle, glass tubing, rubber tubing, pinch clamp, and rubber stoppers. Water is drawn into the apparatus, and the test tube with the reactants is positioned. The system is made airtight, and water is siphoned into a beaker. The test tube is then heated to decompose KClO3 and generate oxygen. The water levels in the apparatus are adjusted, and the volume of water collected is measured. The volume of water collected is equivalent to the volume of oxygen generated.

The calculations in this lab involve determining the weight of oxygen produced from a known weight of KClO3 and Fe2O3. The weight of oxygen is obtained indirectly by measuring the weight of potassium chloride (KCl) left after the oxygen is driven off. The percentage of oxygen in KClO3 is then calculated based on these weights.

The molar gas volume is determined by considering the volume and weight of the collected oxygen. The volume of oxygen at the temperature and pressure in the lab is measured, and the partial pressure of water vapor is subtracted from the barometric pressure to find the pressure of the prepared oxygen. Using the ideal gas law equation (P1V1/T1 = P2V2/T2), the volume of oxygen at standard temperature and pressure (STP) is calculated. From this, the volume occupied by 32 grams (1 mole) of oxygen is determined.

Finally, the density of oxygen is calculated by dividing the mass of oxygen produced by the volume it occupies at STP.

Throughout the lab, students are expected to follow the procedure carefully, make accurate measurements, perform the necessary calculations, and apply the principles of gas laws. The focus is on understanding the relationship between pressure, volume, temperature, and the number of molecules in gases, as well as applying these principles to calculate various quantities related to oxygen.

Objective

·         Use the tools and equipment necessary for basic scientific analysis and research

·         Demonstrate safe practices in the Chemical Laboratory

·         Demonstrate an understanding of gas laws

·         Demonstrate an understanding of Chemical Reaction and Quantities in Chemical Reactions

·         Demonstrate an understanding of Molecules, Compounds and Chemical Composition

·         Demonstrate the proper use of Exponential Notation and Significant Figures

·         Demonstrate an understanding of the composition of matter and energy

·         Record the results of investigation through writing

·         Use scientific reasoning to evaluate physical and natural phenomena

·         Identify the unifying themes of the scientific field of study

Materials

·         8" Test Tube

·         Potassium chlorate (KClO3)

·         Ferric Oxide (Fe2O3)

·         Ring Stand

·         Glass Bottle

·         400 mL Beaker

·         Glass Tubing

·         Rubber Tubing

·         Pinch Clamp

·         Rubber Stoppers

Discussion

Gases at equal temperature and pressure contain an equal number of molecules in their respective volumes. Experimental evidence demonstrates that 32 grams of oxygen (O2) occupies 22.4 liters at STP (0°C and 760 mm Hg). Consequently, any gas occupying 22.4 liters at STP contains the same number of molecules as 32 grams of O2 at STP, and its weight corresponds to the molecular weight of that particular gas. However, in practical experiments, it is not feasible to control the temperature and pressure precisely to STP. Instead, a fraction of the gram-molecule volume is collected at room temperature and atmospheric pressure, which is then converted to STP.

To measure the density of oxygen, a known weight of oxygen is prepared by heating potassium chlorate (KClO3), which easily decomposes to yield potassium chloride and oxygen. The presence of a small amount of Fe2O3 acts as a catalyst. The decrease in weight of the test tube containing KClO3 represents the weight of the generated O2.

This experiment utilizes Dalton's law to determine the pressure of the oxygen gas produced. By applying Boyle's and Charles' laws, the volume of oxygen gas generated can be converted to the corresponding volume at standard temperature and pressure. The fundamental principle underlying this experiment is Avogadro's law.

The measurements obtained from this experiment allow for various calculations, including:

1.       Determining the percentage by weight of oxygen in potassium chlorate.

2.       Measuring the density of oxygen.

3.       Finding the volume of one mole of any gas under standard conditions (the molar gas volume).

Procedure

1.   Place a protective sheet of paper on a weighing scale and measure the weight of an 8" test tube, rounding to the nearest 0.1 mg.

2.   Add 0.8 - 0.9 g of KClO3 to the test tube and weigh it again, rounding to the nearest 0.1 mg.

3.   Repeat the weighing procedure after adding approximately 0.10 8 of Fe203 to the test tube.

4.   Set up the experimental apparatus according to the instructions provided by your instructor (refer to Figure 12).

5.   Ensure that exit Tube A in the bottle reaches the bottom and that the water in the flask is at room temperature.

6.   Draw water through Tube 6 by blowing on it until exit Tube A is filled to the nozzle. Then, tighten the pinch clamp to stop the water from siphoning over.

7.   Place the weighed test tube and its contents in the designated position, ensuring all connections are tight. Open the pinch clamp to allow water to siphon into a small beaker until no more water flows over.

8.   If the system is airtight, the water siphoning will stop almost immediately. While keeping Nozzle D below the water's surface and the clamp open, raise the beaker until the water level in the beaker is equal to that in the flask.

9.   Once the two water levels are aligned, close the pinch clamp and replace the small beaker with a 400 mL beaker accurately weighed to 0.1 g. Open the clamp, allowing a small amount of water to siphon out, which should be retained. Seek approval from your instructor for the apparatus at this stage.

10. With the pinch clamp open, gently heat the test tube to avoid rapid decomposition of KClO3 and excessive pressure buildup in the system.

11. Continue heating until all the oxygen (approximately 250-350 mL of H2O) has been released. Keep Nozzle D below the water level in the beaker and allow Generating Tube B to cool. Some water will be drawn back into the flask.

12. Once Tube B has reached room temperature, raise Beaker E until the water levels in the beaker and flask are once again equal. Close the clamp.

13. Weigh the water in the beaker using a triple beam balance, rounding to the nearest 0.1 gram.

14. Weigh the test tube and record the room temperature and barometric pressure. The volume of water collected is equal to the volume of oxygen generated.

Calculations

(a)     In the initial calculation, we determine the amount of oxygen that can be derived from a measured quantity of potassium chlorate. Heating potassium chlorate leads to its decomposition, resulting in the formation of solid potassium chloride and oxygen gas. Although potassium chlorate decomposes gradually at 400°C, the addition of a catalyst, such as Fe2O3, facilitates rapid decomposition at approximately 270°C.

Since it is challenging to directly weigh gases, the weight of evolved oxygen is determined indirectly through the decomposition of a specific amount of potassium chlorate. This process involves measuring the weight of potassium chloride remaining after the oxygen has been released. By accurately measuring the weight of potassium chlorate before and after the oxygen is driven off, we can calculate the weight of the produced oxygen by subtracting the initial weight from the final weight. Hence, the following conclusion can be drawn:

 Weight of O2 (Wt KClO3 + Wt Fe2O3) - (Wt KCl + Wt Fe203)

These data enable one to calculate the percentage of oxygen in KClO3

% of O2 = ((Weight of O2 ) / (Weight of KClO3)) x 100%

(b)   To determine the molar gas volume, it is necessary to know both the volume and weight of a given amount of oxygen. This experiment enables the measurement of the oxygen volume by displacing water from a bottle. The collected gas volume can be determined by measuring the volume of water displaced. It is important to measure the volume at the temperature and pressure in the laboratory, but the gas laws allow for the calculation of the oxygen volume at standard conditions. Once the volume of a specific weight of oxygen is known at STP, calculating the volume of 32.00 grams (1 mole) becomes straightforward. When correcting to standard conditions, it is essential to account for the fact that the collected gas is saturated with water vapor. The total pressure of the gas is the sum of the partial pressures of oxygen and water vapor.

Subtract the vapor pressure of water used in your experiment from the barometric pressure. This result is the pressure of the prepared oxygen under the conditions at which it was formed.

                                                                           PO2 = Pbar - PH20

It is now possible to calculate the volume the oxygen would occupy under conditions of STP (e.g., 760 mm pressure and 0°C) by use of the following:

Under STP conditions, the volume occupied by the oxygen granules prepared can be determined by measuring the decrease in mass of the test tube. This volume represents the amount of oxygen produced by the quantity of KClO3 used in the experiment. Using this volume and the mass of oxygen obtained, we can calculate the volume that would have been occupied by 32.00 grams (1 mole) of oxygen if it had been prepared.

(c)     The density of oxygen is calculated by division of the Mass of oxygen produced by the Volume that the oxygen occupies at STP.

Report Sheet for THE GRAM MOLECULAR VOLUME of GASES

Name_______________________________________                          Date___________________________

Lab / Section_________________________________

Review Questions

1.       What is the purpose of using a catalyst, such as ferric oxide, in the decomposition of potassium chlorate?

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2.       Explain why equal volumes of gases at the same temperature and pressure contain the same number of molecules.

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3.       Describe the process of collecting oxygen gas by displacing water from a bottle.

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4.       What is Avogadro's law and how does it relate to the experiment?

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5.       How can the weight of oxygen produced be calculated using the weight of potassium chlorate and potassium chloride?

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6.       Why is it necessary to subtract the vapor pressure of water from the barometric pressure in order to determine the pressure of the prepared oxygen?

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7.       Describe the calculations involved in converting the volume of oxygen at room temperature and pressure to the volume at standard temperature and pressure.

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8.       What is the molar gas volume and how is it determined in this experiment?

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9.       Explain the significance of the density of oxygen and how it is calculated in the lab.

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10.    Discuss the relationship between temperature and the volume of a gas, as described by Charles' law.

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11.    How does Dalton's law of partial pressures apply to the determination of the pressure of oxygen gas?

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12.    What are the key steps involved in assembling the experimental apparatus for collecting oxygen gas?

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13.    Explain why it is important to keep the nozzle of the glass tubing below the level of the water in the beaker during the oxygen collection process.

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14.    Describe the factors that can affect the accuracy and precision of the experimental results in this lab.

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15.    Discuss the significance of measuring the weight of the test tube before and after adding potassium chlorate and ferric oxide.

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16.    How does the loss in weight of the test tube indicate the weight of oxygen produced?

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17.    What safety precautions should be followed during the heating of potassium chlorate?

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18.    Explain how the ideal gas law equation (P1V1/T1 = P2V2/T2) is used to calculate the volume of oxygen at standard temperature and pressure.

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19.    Discuss the limitations and potential sources of error in this lab experiment.

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20.    Reflect on the overall objectives of the lab and explain how this experiment contributes to the understanding of gas laws and calculations related to gases.

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