STC_Empirical Formulas of Hydrates Lesson

Empirical Formulas of Hydrates

Determining the empirical formula of a hydrate is easily done using the same method as before. But, instead of determining the formula by each element, you will determine the mole ratio of the compound to water. Look at the example below.

Determine the formula for the hydrate from the given information: 0.391 g Li₂SiF, 0.0903 g H2O Remember the first step of calculating empirical formulas is to convert to moles. ? mol Li SiF = 0.391 g Li SiFx 1 mol Li SiF 155.9 g Li SiF 0.00251 mol Li SiF 1 mol H₂O 18.0 & HO ? mol H,O=0.0903 g H₂O × 0.00502 mol H,O Next, determine the molar ratio by dividing by the smallest. 0.00251/0.00251 = 1 The ratio of water to the salt will become the coefficient. 0.00502/0.00251 = 2 LizSiFe 2H₂O Try this one with percentages: 76.9% CaSO3, 23.1 % H2O ? mol CaSO, = 76.9 g CaSO, × 1 mol CaSO, 120.2 g CaSO, = 0.640 mol CaSO, 1 mol H₂O 18.0 g H,O ? mol H,O= 23.1 gH0 x 1.28 mol H₂O 1.28/0.640 = 2 CaSO3 · 2H2O

Molecular Formulas

When dealing with molecules, the empirical formula may not represent the actual formula of that unknown. For example, you would calculate the empirical formula of hydrogen peroxide to be HO, but the actual formula is H2O2. This whole number ratio of the empirical formula is known as the molecular formula. Sometimes the empirical and molecular formulas are the same, as in the case of water, H2O. If the formula of a molecule can be simplified to smaller whole numbers, then the molecular formula is not the same as the empirical formula.

Which of these compounds are molecular and which are empirical? 

C6H12O6

CO2

NO2

N2O4

C6H12O6 and C2O4 are both molecular because they can be reduced to CH2O and CO2.  CO2 and NO2 are empirical because they cannot be reduced.

Calculating molecular formulas just requires one more step past calculating the empirical formula. Let's review these steps, then add in the new one.

  1. Calculate the # of moles of each element.
  2. Determine the mole ratio (by dividing by the smallest.)
  3. Multiply by integers until you get whole numbers.
  4. Look at the example below to see this last step!

Calculating Molecular Formula Example: Hydrazine has an empirical formula of NH₂. The molecular mass is 32.0 g/mol. What is its molecular formula? First, we must determine the molecular mass of the empirical formula, NH2. 14.01 + (2 x 1.01) = 16.03 g/mol The molecular mass of the molecular formula will always be a multiple of the molecular mass of the empirical formula. To determine that multiple, divide the molecular mass of the molecular formula by the molecular mass of the empirical formula. 32.0 g/mol 16.03 g/mol =2.00 Multiply each subscript in the empirical formula by 2 to get the molecular formula. N(1 x 2.00)H(2 x 2.00) N2H4

Remember to work on the module practice problems as you complete each section of content.

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