Important Differences between Empirical Formula and Molecular Formula

Empirical Formula

An empirical formula represents the simplest, most reduced ratio of elements in a compound. It provides the elemental composition of a substance without detailing the exact arrangement of atoms. This formula is crucial in understanding stoichiometry and chemical reactions. For instance, water has an empirical formula of H2O, denoting that for every two hydrogen (H) atoms, there is one oxygen (O) atom. Empirical formulas are particularly important for compounds with extensive and complex structures, such as polymers or biological molecules. They offer a fundamental insight into the relative proportions of elements in a compound, forming the basis for further study of its molecular structure and properties.

Empirical Formula Example

Example: Ethene (C2H4)

Ethene, also known as ethylene, is a hydrocarbon widely used in chemical industry. Its molecular formula is C2​H4​, which represents the actual number of atoms of each element in a molecule.

To find the empirical formula, we need to determine the simplest whole-number ratio of carbon to hydrogen.

  1. Step 1: Count the Atoms
    • Carbon (C): 2 atoms
    • Hydrogen (H): 4 atoms
  2. Step 2: Find the Ratio
    • The ratio of carbon to hydrogen is 2:4, which can be simplified to 1:2.
  3. Step 3: Write the Empirical Formula
    • The empirical formula for ethene is 2CH

This means that in the simplest, most reduced form, ethene contains one carbon atom for every two hydrogen atoms.

The empirical formula 2CH2​represents the fundamental elemental composition of ethene. It doesn’t provide information about the actual structure or bonding arrangement within the molecule, but it does give us the essential ratio of carbon to hydrogen atoms.

Steps for Empirical Formula Calculation

  • Obtain Composition Data:

Gather information about the elemental composition of the compound. This could be in the form of percentages or actual masses of each element.

  • Convert Masses to Moles:

If given in grams, convert the masses of each element to moles using their respective molar masses (in g/mol).

  • Determine the Smallest Whole Number Ratio:

Divide the number of moles of each element by the smallest number of moles obtained. This will give you the simplest, whole-number ratio of the elements.

  • Write the Empirical Formula:

Using the ratios obtained in step 3, write the empirical formula. It represents the simplest, most reduced ratio of elements in the compound.

  • Optional: Confirm with Experimental Data:

If available, compare the calculated empirical formula with experimental data. They should align if the calculations were accurate.


Let’s say we have a compound with the following composition:

  • 12.01 grams of Carbon (C)
  • 1.01 grams of Hydrogen (H)
  • 16.00 grams of Oxygen (O)
  1. Convert Masses to Moles:
    • Moles of C = 12.01 g / 12.01 g/mol = 1.00 mol
    • Moles of H = 1.01 g / 1.01 g/mol = 1.00 mol
    • Moles of O = 16.00 g / 16.00 g/mol = 1.00 mol
  2. Determine Smallest Whole Number Ratio:
    • The smallest number of moles is 1.00. Divide each by 1.00:
      • C: 1.00 mol / 1.00 mol = 1.00
      • H: 1.00 mol / 1.00 mol = 1.00
      • O: 1.00 mol / 1.00 mol = 1.00
  1. Write the Empirical Formula:
    • The empirical formula is CH2O.

Molecular Formula

The molecular formula represents the exact number of atoms of each element in a molecule. It provides a precise description of the composition, indicating the actual quantities of atoms present. Unlike the empirical formula, which gives the simplest whole-number ratio of elements, the molecular formula offers specific information about the arrangement of atoms within a molecule. For example, glucose has an empirical formula of CH2O, but its molecular formula is C6H12O6, revealing that it contains six carbon (C) atoms, twelve hydrogen (H) atoms, and six oxygen (O) atoms per molecule. The molecular formula is vital in understanding the structure, properties, and behavior of complex compounds in chemistry.

Molecular Formula Calculation

  • Determine the Empirical Formula:

If you don’t already know the empirical formula, you’ll need to calculate it based on the elemental composition.

  • Calculate the Empirical Formula Mass:

Find the sum of the atomic masses of all the elements in the empirical formula. This gives you the molar mass of the empirical formula.

  • Find the Molecular Mass:

Obtain the actual molecular mass of the compound from experimental data or other sources.

  • Calculate the “Factor”:

Divide the molecular mass by the empirical formula mass. This gives you the factor by which the empirical formula needs to be multiplied to get the molecular formula.

  • Multiply the Empirical Formula:

Multiply each subscript in the empirical formula by the factor calculated in the previous step. This gives you the molecular formula.


Let’s consider a compound with an empirical formula of CH2O and a molecular mass of approximately 180 g/mol.

  1. Empirical Formula Mass (CH2O):

    • Carbon (C) = 12.01 g/mol
    • Hydrogen (H) = 1.01 g/mol (x 2 for 2 atoms)
    • Oxygen (O) = 16.00 g/mol
    • Total Empirical Formula Mass = 30.02 g/mol
  2. Factor (F) Calculation:

    • F = Molecular Mass / Empirical Formula Mass
    • F = 180 g/mol / 30.02 g/mol ≈ 6
  3. Molecular Formula:

    • Multiply each subscript in the empirical formula by the factor:
      • C6H12O6

Important Differences between Empirical Formula and Molecular Formula

Basis of Comparison Empirical Formula Molecular Formula
Definition Simplest ratio of elements Exact number of atoms in a molecule
Representation Simple, reduced form Detailed, actual composition
Indicates Minimal elemental composition Specific elemental composition
Multiple Possibilities One empirical formula, multiple molecular formulas possible Only one molecular formula possible
Use Common in covalent compounds Relevant for complex substances
Applicability Applicable to ionic compounds Applicable to covalent and complex substances
Molecular Mass Not used in calculating molecular mass Used to calculate molecular mass
Example H2O (Water) C6H12O6 (Glucose)
Ratio Gives the simplest ratio of atoms Provides the exact count of each atom
Relationship Molecular formula is a multiple of empirical formula Empirical formula is a sub-multiple of molecular formula
Subscripts May or may not be whole numbers Always whole numbers
Common Use Particularly in stoichiometry Widely used in organic chemistry
Example CH2 (Ethene) C6H6 (Benzene)
Variability Can have multiple compounds with same empirical formula Each compound has a unique molecular formula
Molar Mass Empirical formula mass is often smaller Molecular formula mass is typically larger
Composition Represents the simplest whole-number ratio of elements Specifies the actual number of each type of atom

Important Similarities between Empirical Formula and Molecular Formula

  • Elemental Composition:

Both formulas provide information about the types of elements present in a compound.

  • Use in Chemical Formulas:

Both are integral components of chemical formulas, offering insights into the composition of substances.

  • Applicability in Chemistry:

Both formulas play crucial roles in various branches of chemistry, including stoichiometry, organic chemistry, and chemical analysis.

  • Related to Molar Mass:

Both formulas are related to the molar mass of a compound. The molar mass is used to convert between grams and moles.

  • Stoichiometry:

Both formulas are vital in stoichiometric calculations, helping determine the quantitative relationships in chemical reactions.

  • Elemental Ratios:

Both formulas involve ratios of elements, with empirical formulas giving the simplest whole-number ratio and molecular formulas providing the exact count of each type of atom.

  • Representing Chemical Compounds:

Both formulas are used to represent chemical compounds in a concise and standardized manner.

  • Determining Chemical Behavior:

Both formulas contribute to understanding a compound’s behavior in chemical reactions, including its reactivity, stability, and interactions with other substances.

  • Chemical Nomenclature:

Both types of formulas are used in chemical nomenclature to name and represent compounds systematically.

  • Crucial in Organic Chemistry:

Both formulas are fundamental in organic chemistry, where the arrangement of atoms has a significant impact on the properties and reactivity of molecules.

  • Validity Across Compounds:

Both formulas can be applied to a wide range of compounds, from simple substances to complex organic and inorganic compounds.

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