The result of dividing molar mass by empirical mass is a whole number, or a value very close to a whole number. This ratio is used to determine the molecular formula of a compound.
Understanding Molar Mass and Empirical Mass
Before we delve into the result of this division, let's clarify the terms involved:
- Molar Mass: The molar mass of a substance is the mass of one mole of that substance, typically measured in grams per mole (g/mol). It’s found by summing the atomic masses of all atoms in the molecular formula.
- Empirical Mass: The empirical mass is the mass of one mole of the empirical formula of the compound. The empirical formula represents the simplest whole-number ratio of atoms in a compound.
Calculating the Ratio
The provided reference highlights that:
Divide the molar mass of the compound by the empirical formula mass. The result should be a whole number or very close to a whole number.
This whole number result has significance as it helps us determine the molecular formula using the empirical formula.
How It Works
- Determine the Empirical Formula: Obtain the simplest whole-number ratio of atoms in the compound.
- Calculate the Empirical Mass: Find the molar mass of the empirical formula.
- Calculate the Molar Mass: Calculate the actual molar mass of the compound (often given or experimentally determined).
- Divide Molar Mass by Empirical Mass: Divide the molar mass by the empirical mass to get a whole number or a value close to a whole number.
- Determine the Molecular Formula: Multiply all the subscripts in the empirical formula by the whole number obtained in step 4.
Example
Let's say you have a compound with an empirical formula of CH2O and a molar mass of 180 g/mol.
- Empirical Formula: CH2O
- Empirical Mass:
- C: 12.01 g/mol
- H: 2 x 1.01 g/mol = 2.02 g/mol
- O: 16.00 g/mol
- Total Empirical Mass = 12.01 + 2.02 + 16.00 = 30.03 g/mol
- Molar Mass: Given as 180 g/mol
- Divide: 180 g/mol / 30.03 g/mol ≈ 6
- Molecular Formula: Multiply the subscripts in CH2O by 6. Therefore the molecular formula is C6H12O6.
Practical Insights
- This ratio is crucial for determining the molecular formula, which provides the actual number of each type of atom in a molecule.
- The result is practically always a whole number, due to the nature of how molecules are formed and since the empirical formula is the simplest whole number ratio.
- If you get a number far from a whole number, it suggests an error in either the empirical formula, the empirical mass, the molar mass calculation or experimental determination of the molar mass.
Summary
| Step | Description |
|---|---|
| 1. Find Empirical Formula Mass | Calculate the mass of one mole of the empirical formula. |
| 2. Find Molar Mass | Calculate the molar mass of the compound. |
| 3. Divide Molar Mass by Empirical Mass | The result is the factor to determine the molecular formula. |
In essence, dividing the molar mass by the empirical mass yields a whole number that indicates how many times the empirical formula's mass is contained within the molecular formula's mass.
