Gas Laws Exploration

Purpose:To explore the relationships between temperature, pressure and volume of a gas.

Background: The kinetic theory of gases describes some basic assumptions of gases.

  1. Gases are composed of tiny indivisible particles that are in constant, rapid, straight-line motion.
  2. Gas particles possess kinetic energy, due to their motion. The temperature of a gas is the average kinetic energy of its particles.
  3. The collisions between gas molecules are completely elastic.
  4. There is no attraction between particles in a gas.

This has led to the idea of ideal gases, which conform to all assumptions of the kinetic theory. In

reality, ideal gases do not exist. However, the gases that we do encounter deviate only slightly

(often not enough to measure). From these assumptions, several behaviors of gases can be

inferred, and properties explained. In this lab you will explore the relationships between

temperature, pressure and volume of a gas, when one of the three variables is held constant.

Part 1: Charles’ Law: Temperature vs. Volume

Materials:

125 Erlenmeyer flask#4 stopper with 1 holetroughhot plate

thermometer600 ml beakergraduated cylinder

Procedure:

  1. Add 400 ml of water to the 600 ml beaker. Place the stopper in the “dry” flask and then the flask in the beaker of water.
  2. Heat the water to boiling on the hotplate and continue heating for 3-5 minutes. Hold thermometer in boiling water for 2 minutes and record the temperature to the tenths place in the data table. This is the temperature of the gas (air) in the flask and will be considered T1.
  3. Fill a trough with water. Remove the flask from the beaker quickly & carefully. Protect your hand with a paper towel. With one finger over the hole, carefully submerge the flask in the trough upside down so the mouth of the flask is in the water. Caution: Flask is hot! Do not allow any air to enter the flask while transferring to the trough.
  4. Remove finger from stopper under the water. Keep the mouth of the flask under the water at all times. Pour water over the flask to cool it down.
  5. You must equalize the pressure to keep it constant. Raise the flask until the water level inside is equal to the water level outside. The pressure of the gas inside is now equal to the pressure of the gas outside.
  6. Place your finger over the hole in the stopper while the two pressures are equal and quickly invert the flask upright capturing the water that was in the flask. Remove stopper.
  7. Measure the temperature of the water in the flask (air) to the tenths place and record as T2.
  8. Measure the volume of the water remaining in the flask to the tenth’s place with the graduated cylinder and record on line 3 of the data table.
  9. Fill the flask up to the top with water till it’s overflowing. Add the stopper to displace any water over the sink. Carefully remove the stopper and measure the amount of water (air) in the flask to the tenth’s place with the graduated cylinder. This is the total volume of the gas., V1. Record in the data table.
  10. Subtract the volume of air remaining in the flask (step 8) from the total volume of water in the flask (Step 9) to calculate the FINAL VOLUME OF AIR.

Data Table: Experimental Data

Step 2 / Initial temp of hot water (air) T1
Step 7 / Final temp of cold water (air) T2
Step 8 / Volume of water (air) remaining in the flask after the experiment
Step 9 / Total volume of the water (air) in the flask V1
Step 10 / EXPERIMENTAL Final volume of the water (air) V2

Analysis

  1. Write the algebraic equation for Charles’ Law.
  2. Assume the values you have recorded in your data table above (V1, T1, and T2 ) are accurate measurements. Calculate the theoretical V2 . Show Work & express answers in proper significant figures.
  3. Compare your calculated “theoretical” value from and the experimental value for V2seen in the data table. Calculate Percent Error. Show Work & express answer in proper significant Figures.

|Experimental – Theoretical| x 100

|Theoretical|

  1. Which gas variables (moles, volume, pressure or temperature) were held constant during the experiment?

Part 2: Boyle’s Law-Pressure vs. Volume

Procedure:

1.Prepare the sensor by attaching the tubing to the sensor.

2.Move the piston of the syringe to position the front edge of the inside black ring at the 60.0 ml mark and then attach to the tubing of the sensor.

3.Record the pressure of the 60.0 ml volume of contained gas in the data table.

4.Continue moving the piston to the volumes of 52.5 ml, 45.0 ml, 37.5 ml, 30.0 ml, 22.5ml and 15.0 ml, recording the pressure values for each. Make sure you are holding the piston firmly while the value is being recorded.

Data Table

Volume(ml) / Pressure (kPa)

Graph Your Data

  1. Determine the independent variable & dependent variable. Graph the data collected placing the independent variable on the x-axis.
  2. Is the relationship between the variables Inverse or Direct? Explain WHY the pressure changes when the volume changes using the ideas of kinetic molecular theory.

Analysis

  1. If the volume is doubled from 30.0 ml to 60.0 ml, what does your data show happens to the pressure? State your measured pressure values and estimate whether the pressure is doubling, tripling, halving or cutting in thirds.
  2. If the volume is halved from 30.0 ml to 15.0 ml, what does your data show happens to the pressure? State your measured pressure values in and estimate whether it is doubling, tripling, halving or cutting in thirds.
  3. If the volume is tripled from 15.0 ml to 45.0 ml, what does your data show happens to the pressure? State your measured pressure values in and estimate whether it is doubling, tripling, halving or cutting in thirds.
  4. Based on your collected data, what would the pressure value be if you changed the volume of the air in the syringe to 120.0 ml. Use the Boyle’s Law equation to calculate your answer.
  5. Which gas variables (moles, volume, pressure or temperature) were held constant during the experiment?