Chem+Unit+12+Lab

Chem Unit 12 Lab Activity
 * Molar Mass of Butane **

This lab will help us to determine the molar mass of butane, C 4 H 10, experimentally. A simple calculation using a periodic table would give the correct answer for the molar mass of butane, but in this lab, the molar mass of butane will be found with water displacement and the ideal gas law: //PV = nRT.// This formula can be rearranged to solve for //n//, the number of moles: //n = PV/RT//
 * Introduction: **
 * // Butane is a flammable gas, and at NO TIME during this lab should there be any use of an open flame or other heat source! //**

The purpose of the lab is to experimentally determine the molar mass of butane.
 * Objective: **


 * Preparation: **
 * 1) Fill collection tank with room temperature water.
 * 2) Submerge the disposable butane lighter in water, remove it, and then dry it off as thoroughly as possible.
 * Procedure **
 * 1) Weigh the lighter to the nearest 0.01 grams or better.
 * 2) Submerge a 100. ml graduated cylinder in the water so that the cylinder fills completely with water. Invert the cylinder. Make sure there are no air bubbles remaining in the graduated cylinder.
 * 3) Place the lighter underneath the opening of the graduated cylinder in the collection tank. Carefully release the butane from the lighter and collect it in the cylinder. Release enough butane to fill the tube to within 5 ml of its graduated capacity. (95 ml for a 100 ml cylinder). Remove the plastic tube from the graduated cylinder and the lighter.
 * 4) Allow the butane in the cylinder to reach room temperature (about five minutes). Then adjust the level of the water inside and outside the tube to be the same. With the pressure inside the same as the pressure outside, record the volume to the nearest ml. This makes the pressure of the combined water vapor and butane gas equal to the pressure of the atmosphere.
 * 5) Dry off the butane lighter and measure the mass of the lighter again.
 * 6) Record the air temperature in Kelvin.
 * 7) Record the water temperature and barometric pressure.
 * 8) Use the ideal gas law equation to find the molar mass.
 * 9) Use barometric pressure, vapor pressure, water temperature and volume to solve for moles
 * 10) Determine the total mass of butane released into the cylinder by subtracting the two mass measurements.
 * 11) Divide your mass by the number of moles.
 * 12) Repeat the entire procedure 1 more time for a total of 2 trials. Find the average and the % error.

Chem Unit 12 Lab Activity ________/ 50pts
 * Molar Mass of Butane Names: ____________________________________ **


 * Data Table **
 * || Trial 1 || Trial 2 ||
 * Initial mass of lighter ||  ||   ||
 * Final mass of lighter ||  ||   ||
 * Volume of gas collected ||  ||   ||
 * Barometric pressure (given) ||  ||   ||
 * Water temperature ||  ||   ||
 * Vapor pressure of water at room temperature (given) ||  ||   ||
 * Room temperature ||  ||   ||


 * Calculations Table **
 * || Trial 1 || Trial 2 ||
 * Mass of butane released ||  ||   ||
 * Partial pressure of butane ||  ||   ||
 * Moles of butane collected ||  ||   ||
 * Molar mass of butane ||  ||   ||
 * % Error ||  ||   ||


 * Calculations: Please show an example of each calculation in the space provided. **
 * 1) Mass of butane released= Final mass of lighter - Initial mass of lighter.


 * 1) Partial pressure of butane = Barometric pressure – Vapor pressure of water. //(Be sure that pressure units are the same.)//


 * 1) Moles of butane collected = PV/RT. P=partial pressure of butane, T = water temperature in Kelvin.


 * 1) Molar mass of butane = mass of butane released/moles of butane collected.


 * 1) Use a periodic table and determine the accepted value for the molar mass of butane, C 4 H 10 . Then calculate the percent error of both trials.