Carbonation Release Lab Report

Carbonation Release Lab Report Introduction In a chemical reaction where a gas is produced, there will be a pressure change related to the amount of gas that is produced. The Ideal Gas Law states that the pressure, P, (in atm) of n moles of gas occupies a volume, V, (in liters) at temperature T (in Kelvin)1, shown in equation 1. Equation 1. PV = nRT The pressure, P, of a gas and the amount of gas, n, have a directly proportional relationship3. This means that an increase in moles of a gas will result in an increase in pressure when the volume remains constant. This relationship is illustrated in equation 2. Equation 2. P1/n1 = P2/n2 Similarly, pressure and temperature have a directly proportional relationship. When the temperature of a gas is increased, the pressure will increase when the volume remains constant2. Conversely, when the temperature of a gas is decreased, the pressure will decrease when the volume remains constant. This relationship is illustrated in equation 3. Equation 3. P1/V1 = P2/V2 In the carbonation release laboratory experiment, we observed these laws using a thermocouple and a manometer to document the relationships between temperature, pressure, and amount of substance used as a reactant. For this experiment, we used sodium bicarbonate and acetic acid to produce carbon dioxide. We measured the amount of pressure change that resulted from the reaction in equation 4, shown below. Equation 4. NaHCO3(s) + CH3COOH(aq)  CH3COONa(aq) + CO2(g) + H2O(l) Next, we evaluated the data we obtained during this experiment to determine if they were in accordance to what we would expect from the Ideal Gas Law. I hypothesized that our findings will reflect that pressure has a directly proportional relationship with temperature and the amount of gas produced in this experiment. Methods Part I After observing all safety recommendations, we measured at 0.209  0.001g sample of sodium bicarbonate and placed it in a 250mL Erlenmeyer flask. Then we sealed the flask with a rubber stopper with luer lock connections, connected to a manometer to measure the pressure in the flask. We then measured 3.0  0.1mL acetic acid in a syringe, connected the syringe to the luer lock on the stopper of the flask, and began measuring the pressure in the flask. Next we added the acetic acid to the sodium bicarbonate and immediately closed the valve as to not release pressure from the flask. We measured the pressure in the flask at a rate of 10 samples per second and ended the recording after the pressure leveled for several seconds. We repeated this process two more times. For trial 2, we added 0.152  0.001g sodium bicarbonate to 2.2  0.1 mL acetic acid. For trial 3, we added 0.106  0.001g sodium bicarbonate to 1.6  0.1 mL acetic acid. Part II In part two of this experiment, we investigated the effect of temperature on the pressure of a catcher gas. To do this, we prepared a water bath at ambient temperature, which was 23.3°C, or 296.4 5K, in a 1 L beaker. We placed a temperature probe in the beaker and then placed the flask with the gas produced in trial three of the previous section in the water bath. We documented the initial pressure and temperature, and then begin to add ice to the water bath, killing the best slowly and recording both temperature and pressure as we called the gas in the beaker. We also recorded the ambient pressure of the laboratory where we conducted this experiment, at 98.92 kPa. Results With an increase in reactant in this experiment, there was an increase in the amount of gas that was produced. When more gas was produced, there was a higher pressure recorded in the flask. The highest pressure was achieved during trial one, when we use the most sodium bicarbonate and acetic acid. Trial 3, when we used the least moles of reactants, produced the least amount of pressure. Trial 1. Pressure (atm) over time (s) Pressure measured in the fask (kPa) 120 115 110 105 100 95 90 85 0 20 40 60 80 100 120 140 160 180 200 Elapsed time (s) Figure 1. In trial 1, we mixed 17.56 moles sodium bicarbonate with 2.499 x 10-3 moles of acetic acid. This graph shows the pressure in the flask (measured in kPa) over time (measured in s) after the addition of acetic acid to the sodium bicarbonate. Trial 2. Pressure (kPa) over time (s) Pressure measured in the fask (kPa) 112 110 108 106 104 102 100 98 96 94 0 20 40 60 80 100 120 140 160 180 200 Elapsed time (s) Figure 2. In trial 2, we mixed 12.60 moles sodium bicarbonate with 1.833 x 10-3 moles of acetic acid. This graph shows the pressure in the flask (measured in kPa) over time (measured in s) after the addition of acetic acid to the sodium bicarbonate. Pressure measured in the fask (kPa) Trial 3. Pressure (kPa) over time (s) 108 106 104 102 100 98 96 94 0 20 40 60 80 100 120 140 160 180 200 Elapsed time (s) Figure 3. In trial 3, we mixed 8.90 moles sodium bicarbonate with 1.333 x 10-3 moles of acetic acid. This graph shows the pressure in the flask (measured in kPa) over time (measured in s) after the addition of acetic acid to the sodium bicarbonate. In part II of this experiment, we used the product from trial three of part one and cooled the gas in an ice bath from a starting temperature of 23.3° C (296.458K) to 6.6° C (279.7 5K). Our initial pressure in the flask was 106.22 kPa, and our final pressure was 98.25kPa. Gay- Lussac’s Law states that as temperature decreases, pressure will decrease3. We found that to be the case in this part of the experiment. Part 2. pressure measured in the flask as we cooled the flask 108 Pressure (kpa) 106 104 102 100 98 96 94 25 23 21 19 17 15 13 11 9 7 5 Temperature (c) Figure 4. As we cooled the gas contained in the flask, the pressure in the flask decreased, in accordance with Gay-Lussec’s law. Discussion In each trial of part I of this experiment, we decreased the moles of both reactants with each subsequent trial. With a smaller number of moles of reactants, less gas was produced as a product, and therefore there was less resulting pressure in the flask. As shown in figures 1,2, and 3, there is a directly proportional relationship between the moles of reactant used and the moles of gas (and therefore the pressure) produced. Equation 2, shown above, shows the directly proportional relationship between pressure and moles of gas. There are more molecules added to a constant volume, the molecules become more tightly packed, and causes increased pressure within the container2. In part II, we examined the relationship between temperature and pressure of a gas when volume remains constant. As shown in figure 4, pressure has a directly proportional relationship the temperature. As we decreased the temperature of the water bath, the pressure in the flask decreased. As temperature decreases, the energy of the molecules of the gas decreases, causing lower pressure within the flask1. Conclusion With the use of less products in each trailer part one of this experiment, we have pain lower pressure results of products. Highest pressure for trial one was 117.12 kPa, for trial 2 it was 114.8 4 kPa, and for trial 3 it was 104.93 kPa. The results of this part of the experiment were in accordance with the ideal gas law, which states that there is a directly proportional relationship between the moles of reactant and pressure3. Our results supported our hypothesis that more reactant wold result in a higher pressure of the product. In part two, as we decrease the temperature of the iceberg, the pressure within the class decrease. The initial pressure was 106.2 2 kPa at 23.3°C, and her final pressure was 98.25 kPa at 6.6°C, shown in figure 4. These findings are in accordance with Gay-Lussac’s Law, which states that there is a directly proportional relationship between pressure and temperature of a gas when the volume remains constant3. These findings further supported our hypothesis. 1. Jircitano, A. J. http://chemistry.bd.psu.edu/jircitano/gases.html (accessed Nov 9, 2019). 2. Levine, S. Derivation of the Ideal Gas Law. Journal of Chemical Education1985, 62(5), 399. 3. Libretexts. 14.5: Gay-Lussac's Law. https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book:_Introductory_Chemistry _(CK-12)/14:_The_Behavior_of_Gases/14.05:_Gay-Lussac's_Law (accessed Nov 8, 2019). 4. Libretexts. 11.8: Avogadro's Law: Volume and Moles. https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Map %3A_Introductory_Chemistry_(Tro)/11%3A_Gases/11.08%3A_Avogadro’s_Law %3A_Volume_and_Moles (accessed Nov 9, 2019). Appendix Table 1. Target moles of reactants vs. actual reactants used in part I. Trial # Target NaHCO3 Actual NaHCO3 Target CH3COOH Actual CH3COOH (moles) 2.49 x 10-3 (moles) 1 (moles) 2.38 x 10-3 (moles) 2.499 x 10-3 2.50 x 10-3 2 3 1.79 x 10-3 1.19 x 10-3 1.81 x 10-3 1.26 x 10-3 1.92 x 10-3 1.33 x 10-3 1.833 x 10-3 1.333 x 10-3