Introduction:
In this experiment, the molar mas of the chemical used for combustion in a lighter will be determined by using several methods. A flask will be filled to its max capacity and then filled with that gas. After that, the ideal gas law will be used to find mm (the molecular mass of a substance). The equation will be changed to mm=gRT/PV instead of the usual PV=nRT. There are several gas laws exhibited in this experiment. Dalton’s Law of Partial Pressures allows us to determine the pressure of the gas by subtracting the water vapor pressure from the total pressure of the flask. The total pressure would be the barometric pressure for the day and the water vapor pressure would vary with the temperature. The ideal gas law (PV = nRT) is …show more content…
It should be about ¾ full (the exact amount isn’t significant in this experiment).
2) In a 250.0 ml flask, fill it to its max capacity (the very top of the flask). A pipet should be used to ensure accuracy. Record the first volume in the data table that will be used in the first …show more content…
13) Repeat steps 1-12 for the second trial. All data must be recorded in the second data table.
14) Add the two molecular masses together and then divide by two to find the average molecular mass of the C4H10. Also calculate the percent error of the experiment by using the experimental molecular mass of the C4H10 and the actual molecular mass of C4H10 (58g/mol).
Data: First Trial:
Mass of lighter before 11.09 g
Mass of lighter after 10.60 g
Mass of gas (M1 – M2) 0.49 g
Initial V of H2O 308.0 ml
Final V of H2O 99.0 ml
V of gas (V1 – V2) 209.0 ml
T of room 20.9 ° C/ 293.9 K
Barometric Pressure 765.0 torr / 1.01 atm
Water Vapor Pressure at 21° C 18.7 mm Hg /0.025 atm
P of gas (barometric pressure – water vapor pressure) 0.985 atm
Molecular Mass of gas 57.40 g/mol
Second Trial:
Mass of lighter before 10.60 g
Mass of lighter after 10.10 g
Mass of gas (M1 – M2) 0.50 g
Initial V of H2O 307.1 ml
Final V of H2O 87.1 ml
V of gas (V1 – V2) 220.0 ml
T of room 20.9 ° C/ 293.9 K
Barometric Pressure 765.0 torr / 1.01 atm
Water Vapor Pressure at 21° C 18.7 mm Hg /0.025
In this experiment, the molar mas of the chemical used for combustion in a lighter will be determined by using several methods. A flask will be filled to its max capacity and then filled with that gas. After that, the ideal gas law will be used to find mm (the molecular mass of a substance). The equation will be changed to mm=gRT/PV instead of the usual PV=nRT. There are several gas laws exhibited in this experiment. Dalton’s Law of Partial Pressures allows us to determine the pressure of the gas by subtracting the water vapor pressure from the total pressure of the flask. The total pressure would be the barometric pressure for the day and the water vapor pressure would vary with the temperature. The ideal gas law (PV = nRT) is …show more content…
It should be about ¾ full (the exact amount isn’t significant in this experiment).
2) In a 250.0 ml flask, fill it to its max capacity (the very top of the flask). A pipet should be used to ensure accuracy. Record the first volume in the data table that will be used in the first …show more content…
13) Repeat steps 1-12 for the second trial. All data must be recorded in the second data table.
14) Add the two molecular masses together and then divide by two to find the average molecular mass of the C4H10. Also calculate the percent error of the experiment by using the experimental molecular mass of the C4H10 and the actual molecular mass of C4H10 (58g/mol).
Data: First Trial:
Mass of lighter before 11.09 g
Mass of lighter after 10.60 g
Mass of gas (M1 – M2) 0.49 g
Initial V of H2O 308.0 ml
Final V of H2O 99.0 ml
V of gas (V1 – V2) 209.0 ml
T of room 20.9 ° C/ 293.9 K
Barometric Pressure 765.0 torr / 1.01 atm
Water Vapor Pressure at 21° C 18.7 mm Hg /0.025 atm
P of gas (barometric pressure – water vapor pressure) 0.985 atm
Molecular Mass of gas 57.40 g/mol
Second Trial:
Mass of lighter before 10.60 g
Mass of lighter after 10.10 g
Mass of gas (M1 – M2) 0.50 g
Initial V of H2O 307.1 ml
Final V of H2O 87.1 ml
V of gas (V1 – V2) 220.0 ml
T of room 20.9 ° C/ 293.9 K
Barometric Pressure 765.0 torr / 1.01 atm
Water Vapor Pressure at 21° C 18.7 mm Hg /0.025