Molecular Formula: Videos & Practice Problems

Molecular Formula Calculation Example

To determine the molecular formula of a compound, one must first establish its empirical formula and know its molar mass. For instance, consider Ibuprofen, which has a molar mass of 206.3 grams per mole and a percent composition of 75.70% Carbon (C), 8.80% Hydrogen (H), and 15.50% Oxygen (O). The process begins by converting these percentages into grams, assuming a 100-gram sample. This results in 75.70 grams of Carbon, 8.80 grams of Hydrogen, and 15.50 grams of Oxygen.

Next, to find the number of moles of each element, divide the mass of each by its atomic mass, which can be found on the periodic table. For Carbon, the atomic mass is approximately 12 grams per mole, for Hydrogen it is about 1 gram per mole, and for Oxygen, it is 16 grams per mole. This calculation yields approximately 6.3031 moles of Carbon, 8.7302 moles of Hydrogen, and 0.9688 moles of Oxygen.

To simplify these mole values into a ratio, divide each by the smallest number of moles, which in this case is 0.9688. This results in a ratio of approximately 6.5 for Carbon, 9 for Hydrogen, and 1 for Oxygen. Since the ratio must be whole numbers, multiply all ratios by 2 to eliminate the decimal, resulting in 13 Carbons, 18 Hydrogens, and 2 Oxygens. Thus, the empirical formula is C₁₃H₁₈O₂.

The next step is to calculate the empirical mass of the compound using the empirical formula. This involves multiplying the number of each atom by its atomic mass: 13 Carbons (156.13 grams), 18 Hydrogens (18.144 grams), and 2 Oxygens (32 grams). Adding these values gives an empirical mass of approximately 206.274 grams per mole.

To find the n factor, divide the known molar mass of Ibuprofen (206.3 grams per mole) by the empirical mass (206.274 grams per mole). This calculation yields an n factor of approximately 1. Since n equals 1, the molecular formula is the same as the empirical formula, which is C₁₃H₁₈O₂.

In summary, the molecular formula can be derived from the empirical formula by determining the n factor based on the molar mass. In this case, both the empirical and molecular formulas for Ibuprofen are identical, demonstrating that sometimes they can be the same when n equals 1.

Use the given empirical formula and molar mass to determine the molecular formula.
Empirical Formula: NH₂                      Molar Mass: 32.052 g/mol

• A. NH₄
• B. N₄H
• C. NH₂
• D. N₂H₄

Answer: (D)

Glyceraldehyde (M = 90.078 g/mol), a simple monosaccharide, is comprised of 39.999% C, 6.714% H, and 53.297% oxygen by atomic weight. What would be its molecular formula?

• A. C₃H₆O₃
• B. CH₂O
• C. C₃HO₃
• D. C₂H₈O₂

Answer: (A)

Elemental analysis of a pure compound indicated that the compound had 72.2% C, 8.50% H and the remainder as O. If 0.250 moles of the compound weighs 41.55 g, what is the molecular formula of the compound?

• A. C₅H₇O
• B. C₁₀H₁₄O₂
• C. C₆H₈O₂
• D. C₁₂H₁₆O₄

Answer: (B)

Here's what students ask on this topic:

What is the difference between an empirical formula and a molecular formula in chemistry?

The empirical formula represents the simplest whole-number ratio of atoms of each element in a compound. It shows the relative proportions but not the actual number of atoms. In contrast, the molecular formula gives the exact number of atoms of each element in a molecule of the compound. For example, the empirical formula of glucose is $C_1{H}_2{O}_1$ (CH₂O), which shows the ratio of carbon, hydrogen, and oxygen atoms. The molecular formula is $C_6{H}_{12}{O}_6$ (C₆H₁₂O₆), indicating the actual number of atoms in a glucose molecule. Understanding this difference is essential for identifying the composition and properties of chemical compounds.

How do you determine the molecular formula from the empirical formula?

To determine the molecular formula from the empirical formula, you multiply the subscripts in the empirical formula by a whole number called the "n factor." This factor represents how many times the empirical formula units are contained in the molecular formula. For example, if the empirical formula is CH₂O and the n factor is 6, multiplying each subscript by 6 gives the molecular formula C₆H₁₂O₆. The n factor can be found by dividing the molar mass of the compound by the molar mass of the empirical formula. This process ensures that the molecular formula accurately reflects the actual number of atoms in the compound.

What is the "n factor" in molecular formula calculations and how is it used?

The "n factor" is a whole number used to convert an empirical formula into a molecular formula. It represents the multiple of the empirical formula units present in the actual molecule. To use it, multiply each subscript in the empirical formula by the n factor. For example, glucose has an empirical formula CH₂O and an n factor of 6, so its molecular formula is C₆H₁₂O₆. The n factor is calculated by dividing the molar mass of the compound by the molar mass of the empirical formula. This step is crucial when the molecular formula is not directly given and must be derived from experimental data.

How can you find the n factor if it is not given when determining a molecular formula?

If the n factor is not provided, you can calculate it by first determining the empirical formula and its molar mass. Then, obtain the molar mass of the compound through experimental methods such as mass spectrometry. The n factor is found by dividing the compound's molar mass by the empirical formula's molar mass: $n=\frac{\mathrm{MolarMass}}{\mathrm{MolarMassEmpirical}}$. Once you have n, multiply the subscripts in the empirical formula by this number to get the molecular formula. This method ensures accurate identification of the compound's true molecular composition.

Why is it important to understand the relationship between empirical and molecular formulas in chemistry?

Understanding the relationship between empirical and molecular formulas is vital because it helps chemists determine the actual composition of compounds. The empirical formula provides the simplest ratio of elements, which is useful for identifying the basic makeup. However, the molecular formula reveals the exact number of atoms, which is necessary for understanding the compound's properties, behavior, and reactions. This knowledge is essential in fields like pharmaceuticals, materials science, and biochemistry, where precise molecular information guides synthesis and application of chemical substances.