A compound consists of carbon, hydrogen, and fluoride. In one experiment, combustion of 2.50g of the compound produced 3.926 g of CO2. Another sample weighing 5.00 g was found to contain 2.54 g of fluoride. The molar mass is found to be 448.4 g/mol. What are the empirical and molecular formulas? (AW(amu) C=12.01 H=1.008 F=19.00)
I’d like someone to please work through this with me, thank you.
on the CO2, find the mass of C. That is the mass of C in the original, and then the percent of C in the total sample.
for the second sample find the mass of hydrogen (massH=total mass-massF-mass Carbon). For the mass of carbon, use the percent C you found before.
Now you have the percents of the three elements. Assume now, you had 1000 grams, so you can calculate the masses of C, H, F in it. Convert that to moles. Now, you can find the mole ratios which correspond to the subscripts of C, H, and F
Sure! Let's work through this problem step by step.
Step 1: Determine the number of moles of CO2 produced.
First, let's convert the mass of CO2 produced (3.926 g) to moles. To do this, divide the given mass by the molar mass of CO2 (44.01 g/mol):
moles of CO2 = 3.926 g / 44.01 g/mol ≈ 0.0891 mol
Step 2: Calculate the number of moles of carbon in the compound.
Since the combustion of the compound forms CO2, we know that each mole of CO2 contains one mole of carbon. Therefore, the number of moles of carbon in the compound is also 0.0891 mol.
Step 3: Calculate the number of moles of hydrogen in the compound.
The mass of hydrogen in the compound can be calculated by subtracting the mass of carbon (2.50 g) and the mass of fluoride (2.54 g) from the total mass of the compound (5.00 g):
mass of hydrogen = total mass - mass of carbon - mass of fluoride
mass of hydrogen = 5.00 g - 2.50 g - 2.54 g = -0.04 g
Since we can't have a negative mass, it suggests that there is no hydrogen in the compound. Therefore, the number of moles of hydrogen in the compound is 0.
Step 4: Calculate the number of moles of fluoride in the compound.
Similar to the calculation for carbon, we can determine the number of moles of fluoride by dividing the given mass by the molar mass of fluoride (19.00 g/mol):
moles of fluoride = 2.54 g / 19.00 g/mol ≈ 0.134 mol
Step 5: Determine the empirical formula.
The empirical formula represents the simplest whole-number ratio of the elements present in the compound. Since we found that there are 0.0891 mol of carbon and 0.134 mol of fluoride in the compound, the empirical formula would be the ratio of these moles:
Empirical formula = C0.0891F0.134 ≈ CF1.5
Now, we need to multiply all the subscripts by 2 to obtain whole numbers:
Empirical formula = C2F3
Step 6: Calculate the molecular formula.
To find the molecular formula, we need to know the molar mass of the compound. In this case, it is given as 448.4 g/mol. Divide this molar mass by the molar mass of the empirical formula (C2F3) to find the "n" value in the ratio:
n = 448.4 g/mol / (12.01 g/mol × 2 + 19.00 g/mol × 3)
n ≈ 448.4 g/mol / 96.05 g/mol
n ≈ 4.666
Since "n" must be a whole number, we round it to the nearest whole number:
n ≈ 5
Therefore, the molecular formula is:
Molecular formula = (C2F3)n = (C2F3)5 = C10F15
So, the empirical formula of the compound is CF1.5 and the molecular formula is C10F15.