Gas Laws lab report - Gas laws lab PDF

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Gas laws lab...

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Name: Lauren Tanner

Partner: Mary Beazeg

Student No: 20169595

Student No: 20187386

Lab Section: 20

Bench #: 44

Experiment 2: (1 week) (GAS) Gas Laws Purpose In this experiment, the purpose was to explore the relationships between the variables of volume (V), pressure (P), temperature (T), the number of moles (n) and the gas constant (R) = 8.3145 J/mol K). Understanding how these variables relate can be done by learning about the four gas laws, those of which are all used throughout the experiment and on the data collected. This also led to the understanding of the ideal gas law equation, a combination of the four laws.

Introduction After observing a variety of changes in the volume, pressure and temperature of gases, many scientists began to notice how they responded. Their observations of these gases lead these scientists to perform an array of experiments that were very successful and eventually become laws that are used in the present day(4). The first law is Boyle’s Law, (P1)(V1) = (P2)(V2), where volume is inversely proportional to the pressure of an ideal gas when the temperature is constant. The second is Charles’ Law, V1/T1 = V2/T2, which states that volume is directly proportional to the temperature of an ideal gas when it is at a constant pressure. Avogadro’s Law states that the volume of gases have the same number of molecules when both the temperature and pressure are constant, V1/n1 = V2/n2. The final law is Guy-Loussac’s Law, P1/T1 = P2/T2, the pressure is directly proportional to the temperature of an ideal gas when the volume is at a constant. The Ideal Gas Law, PV=nRT was made by combining the four laws into one single equation(1). In theory, an ideal gas would not have a volume or any intermolecular forces acting between the molecules, however, there is no gas that actually behaves like this(2). There are various gases where under specific conditions, can behave almost like an ideal gas. These conditions would be having a high temperature and a low pressure. Nitrogen, Hydrogen and Oxygen are three of the most abundant and common gases in the world, and they behave almost ideally when at room temperature and with a normal atmospheric pressure(3).

Procedure Gather all materials necessary for all three parts of the experiment, which are listed on page 49 of the First Year Laboratory Manual Chemistry 112. Part 1: 1. Set up logger pro on the computer at the lab bench 2. Format it accordingly for volume and pressure 3. Take the syringe and prepare it for data collection 4. Attach the syringe to the pressure sensor, connected to logger pro

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5. Move the plunger of the syringe to the volumes: 5.0 mL, 7.5 mL, 10.0 mL, 12.5 mL, 15.0 mL, 17.5 mL and 20.0 mL 6. Record the pressure at each of the volumes in order to create graphs

Part 2: 1. Format logger pro for pressure and temperature 2. Prepare the rubber flask unit, where the pressure sensor is attached to one of the plastic tips in the stopper 3. Place the Erlenmeyer flask in the boiling water beaker with tongs 4. Record the pressure and temperature, once the readings on logger pro are stable 5. Repeat this process in the hot water, room temperature and ice water beakers 6. Record this data on logger pro to make a graph

Part 3: 1. Weigh 0.10 g of NaHCO3 and place into a 125 mL Erlenmeyer flask 2. Place the rubber stopper into the flask securely 3. From the dispensette, gather 40 mL of acetic acid into a 50 mL beaker and fill the syringe with 1 M of said acid 4. Attach the pressure sensor to one of the plastic tips on the stopper 5. Record the initial temperature of the room temperature bath and the initial pressure of the flask 6. Place the flask into the room temperature beaker and inject the acetic acid into the flask (pull the syringe back to where it was filled to in order to prevent any change in volume) 7. Once the pressure stops rising, record the final temperature and pressure of the bath and flask 8. Use recorded data to create a graph

Questions Part 1

1. Based on your data, what would you expect the pressure to be if the volume of the syringe was increased to 40.0 mL? Explain or show work to support your answer. Based on Boyle’s law that volume is inversely proportional to the pressure of an ideal gas, the pressure would be reduced if the volume of the syringe was increased to 40.o mL.

(P1)(V1) = (P2)( V2) (1.8751 atm)(5.0 mL) = (P2)(40.0 mL) P2 = 0.2344 atm

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2. Find the proportionality constant, k, from the data. If this relationship is direct P=kV, then k = P/V. If it is inverse P=k/V, then k = P×V. Based on your data, Calculate k for both equations and then indicate which formula gives the answer that is “constant” for all your P,V data pairs. Choose the correct one of these formulas and calculate the average k value for the seven ordered pairs in your data table (divide or multiply the P and V values). Calculate is the uncertainty, i.e. standard error, in this final mean value of k? (Use statistics; see the section on determining this in the introduction part of the lab manual) There is an inversely proportional relationship between the variables of pressure and volume, based once again on Boyle’s Law. The value of k can be calculated by using the formula (P)(V) = k. The experimental value of k is 10.3375 and the standard of error is 0.1961, which makes the overall value k = 10.3375 +/- 0.1961 k = P•V =[(9.3755) + (9.84675) + (10.306) + (10.4475) + (10.6155) + (10.77125) + (11.0)] / 7 = 10.3375 Standard Error Using Standard Deviation of Mean:

σ =0.51875 σx=

0.51875 √7

σx=0.1961 Part 2

1. In order to perform this experiment, what two experimental factors were kept constant? The volume (V) in the flask and the amount of the moles (n) were the two experimental factors that were kept constant for this experiment.

2. Write an equation to express the relationship between pressure and temperature (K). Use the symbols P, T, and k. Based on the ideal gas equation, PV=nRT, the variables volume (V), R (gas constant) and number of moles (n) can be represented as k, due to them being the constants within the experiment. This leaves pressure and temperature which has a direct relationship, leading to the equation k=P/T.

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3. Find the proportionality constant k. Like above, if the relationship is direct, k = P/T. If it is inverse, k = PxT. Based on your answer to Question 2, choose one of these formulas and calculate k for the four ordered pairs in your data table (divide or multiply the P and T values). Show the answer in the fourth column of the Part 2 data table. How “constant” were your values?

Using the equation k = P/T, the calculated values of k shown in the data table below express values that are not consistent. They do not signify a constant value as it continually decreases as the temperature and pressure continually decline. This error that occurred during the experiment could be due to any number of factors including a leak in the flask’s seal, with the temperature probe or Logger Pro.

4. The data that you have collected can also be used to determine the value for absolute zero on the Celsius temperature scale. On the plot of Celsius temperature on the y-axis and pressure on the x-axis, find absolute zero. Y=mx+b 0 = (0.0056)x + 0.4909 -0.4909 = 0.0056x X = -87.66

5. Since “absolute zero” is the temperature at which the pressure theoretically becomes equal to zero, the temperature where the regression line (the extension of the temperature-pressure curve) intersects the y-axis (b in the y=mx+b equation) should be the Celsius temperature value for absolute zero. What is your experimental value for absolute zero? What is the uncertainty (do a visual estimate based on your plot)? Is your answer correct, i.e. does the real value lie within the limits of your experimental uncertainty? The experimental value for absolute zero that was calculates was -87.66 +/- 0.1, which converts to 183.49 K. This does not match the theoretical value of absolute zero. The experimental value does not fall into the range of uncertainty however, this could be causes by any limitations that arose during the experiment as the laboratory was unable to calculate values to an infinite value.

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Part 3

1. Complete the data sheet and show all calculations in your notebook.

Initial Amount of Gas: PV = nRT

(1.0086 atm)(155) = n(8.3145)(296.45) n = (1.0086 atm)(155) / (8.3145)(296.45) n = 0.0634 mol Final Amount of Gas: PV = nRT (1.1581)(155) = n(8.3145)(296.05) n = (1.1581)(155) / (8.3145)(296.05) n = 0.0728 mol Amount of CO2 added: CO2 = Final – Initial = 0.0728 – 0.0634 =9.4 x 10^3 mol 2. Use Dalton’s law to determine the partial pressure of the CO2 formed. P total = P1 + P2 1.1581 = 1.0086 + P2 P2 = 1.1581 – 1.0086 = 0.1495 atm

3. Based on the known values: partial pressure (pressure change seen during the reaction), the volume of the flask, the number of moles of CO2 (found using mass of sample and molar mass of NaHCO3) and the final temperature in Kelvin (conversion: T in Celsius + 273.15) what is the value of the gas constant R? – Don’t forget units. Use PV=nRT and solve for R. Determine the uncertainty in R.

PV = nRT (0.1495)(155) = (9.4 x 10^3) R (296.05) R= (0.1495)(155) / (9.4 x 10^3) (296.05) R= 8.3268 J/K * mol

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4.Compare this value to the accepted value. (8.3145 J K -1 mol-1). In this context, compare means that you should determine if the literature value for R agrees with your experimental value within your limits of uncertainty. The calculated value of R, as shown above, is 8.3268 J/K * mol, which almost agrees with the accepted value of R = 8.3145 J/K * mol. The R value calculated is only off of the accepted value by 0.0123 which makes it fairly accurate. Although the value calculated does agree for the most part, there are some possible reasons as to why it is not even closer. This could be due to limitations that come from the equipment used in this experiment as it may not be able to determine the exact infinite value of R. In the Ideal Gas Equation, R is calculated using the variables of an ideal gas, however, these properties do not correspond with any known gas, making it difficult to achieve such a result from the data collected during this experiment.

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DATA SHEET Part 1 Volume (mL)

Pressure (atm)

Constant, k

P•V

5

1.8751

P/V

0.37502 9.3755

7.5

1.3129

9.84675

0.175053

10

1.0306

10.306

0.10306

12.5

0.8358

10.4475

0.066864

15

0.7077

10.6155

0.04718

17.5

0.6155

10.77125

0.03517

20

0.5500

11

0.0275

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Water Bath

Pressure (atm)

Temperature (°C)

Temperature (K)

Constant, k (P / T or P•T)

1.0671

100

373.15

2.86 x 10^-3

0.6885

40.5

313.65

2.12 x 10^-3

0.6172

23.9

297.05

2.07 x 10^-3

0.5581

7.2

280.35

1.99 x 10^-3

Boiling

Hot Room temp Ice

Part 2 Part 3

Mass of NaHCO3 0.10 g Volume of flask 155 mL Temperature (initial) 23.3 ˚C Temperature (final) 22.9 ˚C Pressure (initial) 1.0086 atm Pressure (final) 1.1581 atm Amount of gas initially in flask

0.063 mol

Amount of CO2 added 9.4 x 10^3 mol

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References [Each reference listed here should have a number that corresponds to the superscripted number in the body of the report where this reference is pertinent.] 1. “Experiment 2 -Gas Laws (47-48). Queen’s Chemistry First Year Laboratory Manual Chemistry 112, 2019-2020 edition, Department of Chemistry Ideal Gas Law. (n.d.). Retrieved October 2, 2019, from 2. https://www.sciencedirect.com/topics/engineering/ideal-gas-law. 3. Helmenstine, A. M. (2018, June 16). What Are the 4 Most Abundant Gases in Earth's Atmosphere? Retrieved October 2, 2019, from https://www.thoughtco.com/most-abundantgases-in-earths-atmosphere-607594. 4. Gas Laws. (n.d.). Retrieved October 2, 2019, from http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch4/gaslaws3.html....