if 16.5 of iron (II) sulfide contain 6.02 x 10^23 formula units, how many formula units of iron (II) sulfide are in 1.00 mL of iron (II) sulfide? The density of iron (II) sulfide is 4.84 g/cm^3
I can tell you without any work at all that 16.5 grams of FeS does not contain 6.02 x 10^23 formula units of FeS.
To find out how many formula units of iron (II) sulfide are in 1.00 mL, follow these steps:
Step 1: Calculate the mass of iron (II) sulfide in 1.00 mL using the given density.
Density = Mass / Volume
Rearrange the formula to solve for Mass: Mass = Density x Volume
Mass = 4.84 g/cm^3 x 1.00 mL
Note that the units cancel out, leaving you with grams.
Step 2: Convert the mass of iron (II) sulfide to formula units.
To do this, you need to know the molar mass of iron (II) sulfide. The formula for iron (II) sulfide is FeS.
The atomic masses of Iron (Fe) and Sulfur (S) are:
Fe = 55.85 g/mol
S = 32.07 g/mol
So, the molar mass of FeS = 55.85 g/mol + 32.07 g/mol = 87.92 g/mol
Now, divide the mass of iron (II) sulfide in grams (from Step 1) by the molar mass to get the number of moles:
Moles of FeS = Mass of FeS / Molar mass of FeS
Step 3: Convert moles of iron (II) sulfide to formula units.
Avogadro's number states that one mole of any substance contains 6.02 x 10^23 particles (formula units in this case). So, multiply the number of moles (from Step 2) by Avogadro's number to get the number of formula units:
Formula units of FeS = Moles of FeS x Avogadro's number
Solving these steps using the given values:
Step 1: Mass = 4.84 g/cm^3 x 1.00 mL = 4.84 g
Step 2: Moles of FeS = 4.84 g / 87.92 g/mol ≈ 0.055 moles
Step 3: Formula units of FeS = 0.055 moles x 6.02 x 10^23 formula units/mole ≈ 3.31 x 10^22 formula units
So, there are approximately 3.31 x 10^22 formula units of iron (II) sulfide in 1.00 mL.