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1.4.3Additional lessons

Empirical Formula — IB Chemistry HL Revision Notes & Topical Questions

5 min read 5 parts Lesson 16 of 172

📚 IB Chemistry HL Hub / Topical Revision
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Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
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Section 1 of 5Types of Formula

Different formulae describe a compound in different levels of detail.

Empirical formula
  • Empirical formula The simplest whole-number ratio of atoms of each element in a compound.
Molecular formula
  • Molecular formula The actual number of atoms of each element in one molecule, which is a whole-number multiple of the empirical formula.
Example
  • Example Glucose has molecular formula C₆H₁₂O₆ and empirical formula CH₂O.
Ionic compounds
  • Ionic compounds The formula unit is always the empirical formula because no discrete molecules exist.

Section 2 of 5Finding the Empirical Formula

Use masses or percentages, then follow these steps.

1
Write the mass of each element, treating the percentage as grams in a 100 g sample.
2
Divide each mass by the relative atomic mass to obtain moles of atoms.
3
Divide all the mole values by the smallest to find the ratio.
4
If the ratio is not whole numbers, such as 1.5, multiply all by a suitable factor to get whole numbers.
5
Write the empirical formula using these whole-number ratios as subscripts.

Section 3 of 5Finding the Molecular Formula

Finding the Molecular Formula
Mass of empirical unit
  • Mass of empirical unit Add the atomic masses of the atoms in the empirical formula.
Ratio
  • Ratio Divide the given molar mass by the empirical formula mass to find a whole number n.
Multiply
  • Multiply Multiply every subscript in the empirical formula by n.
Source of M
  • Source of M The molar mass is found from mass spectrometry or another measurement.

Section 4 of 5Worked Example

Question

  • A compound is 40.0% C, 6.7% H and 53.3% O by mass, with M = 180 g mol⁻¹.
  • Find its empirical and molecular formulae.

  1. Moles
    • C = 40.0 ÷ 12.01 = 3.33;
      • H = 6.7 ÷ 1.01 = 6.63;
      • O = 53.3 ÷ 16.00 = 3.33.
  2. Divide by 3.33: C : H : O = 1 : 2 : 1, so the empirical formula is CH₂O (M = 30.03).
  3. 180 ÷ 30.03 ≈ 6, so the molecular formula is C₆H₁₂O₆.

Answer

  • Empirical formula CH₂O;
    • molecular formula C₆H₁₂O₆

Section 5 of 5Exam Checklist & Tips

1.4.3 – Key Points for the Exam

  • Empirical formula The simplest whole-number ratio of atoms in a compound.
  • Molecular formula The actual numbers of atoms of each element in a molecule.
  • Percentage composition The percentage by mass of each element in a compound.
  • Mole ratio The ratio of amounts of different elements or species.

Common Mistakes

  • Rounding a ratio such as 1.5 down to 1 instead of multiplying to a whole number.
  • Using the mass numbers rather than atomic masses.
  • Dividing by the largest number of moles instead of the smallest.
  • Giving the empirical formula when the molecular formula is requested.

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