Empirical Formulas vs Molecular Formulas - Explained

Introduction

This tutorial will help you understand the difference between empirical formulas and molecular formulas in chemistry. By the end of this guide, you will be able to distinguish between the two types of formulas, calculate them from given data, and apply this knowledge to solve problems effectively.

Step 1: Understanding Empirical Formulas

Empirical formulas represent the simplest whole-number ratio of atoms in a compound.

• Definition: An empirical formula shows the relative number of each type of atom in a compound, not the actual number of atoms.
• Example: The empirical formula for glucose (C6H12O6) is CH2O, as it simplifies to the smallest ratio of carbon, hydrogen, and oxygen atoms.

Practical Advice

• To find the empirical formula:
  1. Determine the mass of each element present in the compound.
  2. Convert these masses to moles by dividing by the atomic mass of each element.
  3. Divide the number of moles of each element by the smallest number of moles calculated.
  4. Round to the nearest whole number to get the ratio.

Step 2: Understanding Molecular Formulas

Molecular formulas indicate the actual number of each type of atom in a molecule.

• Definition: A molecular formula is a multiple of the empirical formula, which provides the actual quantities of atoms.
• Example: For glucose, its molecular formula is C6H12O6, indicating there are six carbon atoms, twelve hydrogen atoms, and six oxygen atoms.

Practical Advice

• To find the molecular formula:
  1. First, determine the empirical formula as outlined in Step 1.
  2. Calculate the molar mass of the empirical formula.
  3. Divide the molar mass of the compound by the molar mass of the empirical formula.
  4. Multiply the subscripts in the empirical formula by this whole number to obtain the molecular formula.

Step 3: Working Through Examples

To solidify your understanding, let’s work through an example.

Example 1: Finding an Empirical Formula

• Given: A compound contains 40% sulfur and 60% oxygen by mass.
  1. Convert the percentages to grams (assume 100g total).
  2. Calculate moles:
     • Sulfur: 40 g / 32.07 g/mol ≈ 1.25 moles.
     • Oxygen: 60 g / 16.00 g/mol ≈ 3.75 moles.
  3. Divide by the smallest number of moles (1.25):
     • Sulfur: 1.25 / 1.25 = 1
     • Oxygen: 3.75 / 1.25 = 3
  4. The empirical formula is SO3.

Example 2: Finding a Molecular Formula

• Given: The empirical formula of a compound is CH2 and its molar mass is 42 g/mol.
  1. Calculate the molar mass of the empirical formula:
     • C: 12.01 g/mol, H: 2.02 g/mol = 14.03 g/mol.
  2. Divide the compound's molar mass by the empirical formula’s molar mass:
     • 42 g/mol / 14.03 g/mol ≈ 3.
  3. Multiply the subscripts in the empirical formula by 3:
     • C3H6.
  4. The molecular formula is C3H6.

Step 4: Practice Problems

To reinforce your learning, try solving the following problems:

• Calculate the empirical formula for a compound with 50% carbon and 50% hydrogen.
• Determine the molecular formula of a compound with an empirical formula of C2H5 and a molar mass of 70 g/mol.

Conclusion

In this tutorial, you learned the essential differences between empirical and molecular formulas and how to calculate each. Understanding these formulas is crucial in chemistry as they provide insights into the composition of compounds. Practice the example problems provided to enhance your skills, and consider exploring more complex compounds as your confidence grows.