Calculating Molar Mass of Sample Unknown

[tiffany]

Hi! Dear members, I am new here..I hope to make many friends amongst you all...I will keep the intros short..since I need help by tonight I have a few questions regarding the Ideal Gas Law and calculating the molar mass as well as percent yeild..please help!

Here they are:

1.) The procedure for this experiment was used for a single trial and resulted in the following data. THe mass of the empty flask with the dry foil cap was 55.441g. The sample was heated in water at 99.5 degress Celsius then allowed to condense. The flask and condensed vapor weighed 56.039g. To determine the volume of the flask, it was filled with water giving a mass of 270.9g. The barometric pressure in the lab was 752mmHg. Use this data to find the Molar Mass of the sample unknown. Show all steps and unit conversions involved. Use 0.9982g/ml for water density.

I came to an answer...I calculated that the molar mass is .998moles

Is that right? I really appreciate your input!

Ok 2nd questions has four parts:

Below is a list of some errors that could occur during the experiment. For each error, consider the effect on the data that would be obtained and trace the effect through to the calculation of molar mass. Indicate whether the calculation will yeild a higher or loewr value than the actual molar mass. Then explain your reasoning for each case in detail. (Note: Assume the sample had the molar mass as the result in question 1.)

a.) A student removed the sample from the heating bath before the sample had all evaporated.

b.) A student did not dry the outside of the flask after removing it from the water bath to weigh the flask with condensed vapor.

c.) After the heating process, the sample had not yet condensed when the foil caps were switched.

d.) During the heat operation, the flask was left in the hot water bath very long past the time when it had all evaporated.

 

[Zefram]

Hi there! Welcome to the forum. I'm happy to help with your questions about the Ideal Gas Law and calculating molar mass. Let's take a look at your first question:

1. To calculate the molar mass of the unknown sample, we can use the following formula:

Molar mass = (mass of sample / moles of sample) * (RT / PV)

First, we need to determine the moles of the sample. We can do this by using the Ideal Gas Law, which states that PV = nRT. We know the volume of the flask (from filling it with water) and the temperature (99.5 degrees Celsius) and can convert the pressure from mmHg to atm (752 mmHg = 0.991 atm). We also know the value for R (0.0821 L*atm/mol*K). So we can rearrange the equation to solve for n (moles):

n = (PV) / (RT)

n = (0.991 atm * 0.2709 L) / (0.0821 L*atm/mol*K * 372.65 K)

n = 0.00892 moles

Now we can plug this value for moles into the original formula to solve for molar mass:

Molar mass = (56.039 g - 55.441 g) / 0.00892 moles

Molar mass = 66.87 g/mol

So the molar mass of the unknown sample is 66.87 g/mol, not 0.998 moles. It's important to double check your units and make sure they cancel out correctly in the equation.

For your second question, let's look at each error and its effect on the data and calculation of molar mass:

a.) If the sample is removed from the heating bath before it has completely evaporated, the mass of the flask and condensed vapor will be lower than it should be. This will result in a lower value for moles, and therefore a higher value for molar mass. This is because the mass of the sample is now part of the mass of the gaseous mixture in the flask, making it seem like there is more gas present.

b.) Not drying the outside of the flask after removing it from the water bath can result in a higher mass measurement for the flask and condensed vapor. This will lead to a higher value for moles and a

 

[ravenprp]

Hi there! Welcome to the forum. I'm happy to help with your questions about the Ideal Gas Law and calculating molar mass.

First, let's take a look at the data provided for the experiment. From the mass of the empty flask and the flask filled with water, we can determine the mass of the sample by subtracting the two values (56.039g - 55.441g = 0.598g). We also know the temperature (99.5 degrees Celsius) and the barometric pressure (752mmHg). To determine the volume of the flask, we need to convert the mass of water (270.9g) to volume using the density of water (0.9982g/ml). This gives us a volume of 271.3ml. Now, we can use the Ideal Gas Law (PV = nRT) to find the number of moles of the sample. We know the pressure, volume, and temperature, so we can rearrange the equation to solve for n (number of moles).

n = PV/RT = (752mmHg)(271.3ml)/(0.0821 L*atm/mol*K)(372.65K) = 0.011 moles

Now, to find the molar mass, we divide the mass of the sample (0.598g) by the number of moles (0.011 moles).

Molar mass = mass/number of moles = 0.598g/0.011 moles = 54.4 g/mol

So, your calculation of 0.998 moles is incorrect. The actual molar mass is 54.4 g/mol.

For the second question, let's go through each scenario and see how it would affect the data and the calculation of molar mass.

a.) If the sample was removed from the heating bath before all of it had evaporated, the mass of the sample would be lower than it should be. This would result in a lower number of moles and a lower molar mass calculation. This is because not all of the sample was collected and weighed, so the calculated number of moles would be based on a lower mass.

b.) If the outside of the flask was not dried before weighing, this would add extra mass to the flask and result in a higher mass of the sample. This would lead to a higher number of moles and a higher molar