Title: Gas Laws: Boyle’s Law, Estimating Absolute Zero
Date: March 13, 2015
Report Writer: John Torri
Lab Partners: Steven, Justin
Abstract
In this experiment we tested and identified the different components that make
up a gas. Including Volume, Temperature, Pressure, and Moles. Using different
experiments and measurement equipment we were able to identify the components of
the gas. With this information we can determine the ideal gas constant.
Materials

Lab Quest tool

Beaker

Mg Ribbon
Procedure:
Part 1.
The goal of Part 1 was to determine the relationship between the pressure and
the volume of a confined gas when the temperature is constant, and to experimentally
estimate the temperature of absolute zero. Using a syringe connected to a LabQuest
pressure sensor we measured the gasses pressure at 7 different points and recorded
them in our post lab. Based on the data gathered we saw the correlation between the
pressure at different volumes. As the volumes got larger the pressure decreased
confirming Boyle’s Law.
Pressure (torr)
Table 1. Relationship between
Volumes and Pressure of a Gas
1600
1400
1200
1000
800
600
400
200
0
Volume
(ml)
10
5
7.5
12.5
15
17.5
20
Volume (ml)
Figure 1.1) Table 1. Displays the correlation of different pressures (torr) at different
volumes (ml).
Pressure (torr)
Table 2. Relationship between
Pressure and Volume
1600
1400
1200
1000
800
600
400
200
0
20
17.5
15
Volume (ml)
12.5
10
Figure 1.2) Table 2. Displays a reverse measurement of the change in volume (ml) and
the pressure change (torr).
Part 3.
The goal of part three was to measure the value of the Ideal Gas Constant.
Formula.
Using a strip of Mg ribbon (3cm) we added 3M HCl solution (20ml) to the Graduated
cylinder until it reached the 40ml mark. Once the HCL solution is combined with the
water we then added the Mg strip in the tube then added the stopper to close off the
tube. Immersing the tube in the graduated cylinder. Within a minute the Mg ribbon
reacts and disappears in the solution.
We recorded the mass of the Mg Ribbon, the water temperature of the
Graduated cylinder, Volume, and determined the pressure of the reaction. Using the
data we gathered we were then able to calculate gas constant, R in L atm moleK. Our
results told us that R= .0894atmL/moleK.
Part 4.
In Part four the objective was to investigate the relative rates at which two gases
of different molar masses diffuse, and to compare the observed relative rates with that
predicted by Graham’s Law.
Using a 24 inch tube, and a Q-tip (broken in half), and two small one hole
stoppers. Each half of the Q-tip was dipped in either an acid (HCl)or a base (NH4OH).
Then simultaneously placed on either sides of the 24 inch tube, we waited until a faint
white line formed on the tube. Once the line formed we measured the distance from
the tip of the Q-tip on each end. Using this ratio we plugged it to Grahams formula. We
divided the distance of MM(HCl)/MM(NH4OH) and received the value of 1.24cm.
Calculations:
Part 1. (Boyle’s Law)
Volume (ml)
Pressure (torr)
10
768.7
5
1490.1
7.5
1004.1
12.5
618.7
15
517.7
17.5
444.8
20
391.3
Volume (mL)
20
17.5
15
12.5
10
Pressure (torr)
767.3
873.1
1015.5
1208.6
1492.7
Part 3. (Ideal Gas Law)

Mg Mass: .0035g -> 1mole Mg/ 29.31g= .0014M Mg

Volume= 36.9mL/1000L = .0369L

Pressure= 755.3torr-> 1 atm/760torr= .994 atm

Temp= 21 C (21+273= 294K)


P= (.994 atm) x V= (.0369L)/ (M=.0014) x (T=294K)

.0894 atmL/Mole K
Part 4. (Graham’s Law)
Distance of HCl: 16.9cm
Distance of NH4OH: 21.0cm
Ratio: 21cm(NH4OH)/ 16.9cm (HCl) = 1.24cm
Conclusions:
This lab solidified and supported the laws and rules we have learned
throughout our classes. Using Boyle’s Law in we saw in this experiment how Volume
and pressure were directly related and how the larger volume objects had a lower
amount of pressure. Part three of the lab showed the ideal gas law and how the
different components of V,P, n, T can be used in the Ideal Gas formula to find the
missing variable. Part four dealt with Grahams law, measuring the distances of each
gas and determining the square root of their molar masses, both at the same temp,
and pressure. In conclusion, all of these law’s and formulas helped to identify these
gasses and determine other information in relation to each formula.