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Molecular Weight Calculator

Chemical Formula, Atomic Mass Calculator • 2026

Molecular Weight Formulas:

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\( \text{Molecular Weight} = \sum (\text{Atomic Mass} \times \text{Number of Atoms}) \)

\( \text{Mass Percent} = \frac{\text{Mass of Element}}{\text{Molecular Weight}} \times 100 \)

\( \text{Moles} = \frac{\text{Mass (g)}}{\text{Molecular Weight (g/mol)}} \)

\( \text{Mass (g)} = \text{Moles} \times \text{Molecular Weight (g/mol)} \)

\( \text{Empirical Formula} = \text{Simplest whole number ratio of atoms} \)

Where:

  • \( \text{Molecular Weight} \) = Sum of atomic masses in a molecule
  • \( \text{Atomic Mass} \) = Average mass of atoms of an element
  • \( \text{Number of Atoms} \) = Subscript in chemical formula
  • \( \text{Mass Percent} \) = Percentage by mass of an element
  • \( \text{Moles} \) = Amount of substance
  • \( \text{Empirical Formula} \) = Lowest ratio of elements

Molecular weight calculations are fundamental in chemistry, biochemistry, and pharmaceutical applications. These formulas help scientists determine stoichiometric relationships, prepare solutions, and analyze chemical compositions.

Example: For H₂O: MW = (2 × 1.008) + (1 × 15.999) = 2.016 + 15.999 = 18.015 g/mol.

Chemical Formula

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Results

18.015 g/mol
Molecular Weight
0.555 mol
Calculated Moles
10.00 g
Calculated Mass
3
Total Atoms
H₂O
Formula: H₂O
MW: 18.015 g/mol
Key Formulas
Molecular Weight
MW = Σ(atomic_mass × count)
MW = (A₁×n₁) + (A₂×n₂) + ...
Mass Percent
% = (mass_element / MW) × 100
n = mass / MW
Mole Calculations
moles = mass / MW
mass = moles × MW
Molecular Statistics
Formula
H₂O
Molecular Weight
18.015 g/mol
Total Atoms
3
Elements
H, O
Element Count Atomic Mass Contribution
Mass Parameter Value Unit Formula
Mole Parameter Value Unit Description

Molecular Weight Fundamentals

What is Molecular Weight?

Molecular weight (also called molecular mass) is the sum of the atomic weights of all atoms in a molecule. It's expressed in atomic mass units (amu) or grams per mole (g/mol). This fundamental property is essential for stoichiometric calculations, solution preparation, and chemical analysis.

Key Chemistry Metrics

Molecular Weight (MW): Sum of atomic masses in a molecule

Atomic Mass: Average mass of atoms of an element

Formula Weight: Molecular weight of ionic compounds

Mass Percent: Percentage by mass of each element

Empirical Formula: Simplest whole number ratio of atoms

Key Rules:
  • MW = Σ(atomic_mass × atom_count)
  • Subscripts indicate atom counts
  • Mass percent = (element_mass / MW) × 100
  • Moles = mass / MW
  • Mass = moles × MW
  • Use periodic table for atomic masses

Applications

Practical Applications

Molecular weight calculations are essential for preparing standard solutions, performing stoichiometric calculations, analyzing chemical compositions, and conducting quantitative analytical chemistry. They're fundamental in laboratories, pharmaceutical preparation, and biochemical research.

Practical Uses
  1. Standard solution preparation
  2. Stoichiometric calculations
  3. Chemical composition analysis
  4. Pharmaceutical formulations
  5. Quality control testing
  6. Biochemical assays
  7. Reaction yield calculations
Laboratory Guidelines:
  • Use accurate atomic masses from periodic table
  • Account for hydrates and solvents
  • Consider isotopic abundance
  • Verify chemical formulas
  • Round appropriately for significant figures

Molecular Weight Learning Quiz

Question 1: Multiple Choice - Basic Molecular Weight

What is the molecular weight of glucose (C₆H₁₂O₆)? (Atomic masses: C=12.01, H=1.008, O=16.00)

Solution:

The answer is A) 180.16 g/mol. Calculate: (6 × 12.01) + (12 × 1.008) + (6 × 16.00) = 72.06 + 12.096 + 96.00 = 180.156 ≈ 180.16 g/mol.

Pedagogical Explanation:

This question tests the fundamental molecular weight calculation. Each subscript in the formula indicates the number of atoms of that element. The molecular weight is the sum of all atomic contributions. This calculation is essential for stoichiometric relationships.

Key Definitions:

Molecular Weight: Sum of atomic masses in a molecule

Atomic Mass: Average mass of atoms of an element

Subscript: Number of atoms in the formula

Important Rules:

• MW = Σ(atomic_mass × atom_count)

• Subscripts indicate atom counts

• Use accurate atomic masses

Tips & Tricks:

• Organize by element when calculating

• Double-check subscripts

• Use periodic table for atomic masses

Common Mistakes:

• Forgetting to multiply by subscripts

• Using incorrect atomic masses

• Misreading subscripts

Question 2: Detailed Answer - Mass Percent Calculation

Calculate the mass percent of carbon in ethanol (C₂H₅OH). (Atomic masses: C=12.01, H=1.008, O=16.00)

Solution:

Step 1: Calculate molecular weight of C₂H₅OH

Carbon: 2 × 12.01 = 24.02 g/mol

Hydrogen: 6 × 1.008 = 6.048 g/mol

Oxygen: 1 × 16.00 = 16.00 g/mol

MW = 24.02 + 6.048 + 16.00 = 46.068 g/mol

Step 2: Calculate mass percent of carbon

Mass % of C = (Mass of carbon / MW) × 100

Mass % of C = (24.02 / 46.068) × 100 = 52.14%

Therefore, carbon makes up 52.14% of ethanol by mass.

Pedagogical Explanation:

This problem demonstrates mass percent calculation, which is important for composition analysis. Mass percent shows the contribution of each element to the total molecular weight. This information is crucial for analytical chemistry and material characterization.

Key Definitions:

Mass Percent: Percentage by mass of an element in a compound

Composition Analysis: Determining elemental makeup

Quantitative Analysis: Measuring amounts of substances

Important Rules:

• Mass % = (element_mass / MW) × 100

• Sum of all mass % = 100%

• Essential for analytical chemistry

Tips & Tricks:

• Always verify sum of mass percents = 100%

• Use for empirical formula determination

• Critical for purity analysis

Common Mistakes:

• Forgetting to multiply by 100 for percentage

• Using incorrect molecular weight

• Not accounting for all atoms of element

Question 3: Word Problem - Mole Calculation

A chemist needs 0.25 moles of sodium chloride (NaCl) for an experiment. How many grams of NaCl should be weighed? (Atomic masses: Na=22.99, Cl=35.45)

Solution:

Step 1: Calculate molecular weight of NaCl

MW = 22.99 + 35.45 = 58.44 g/mol

Step 2: Calculate mass needed

Mass = moles × MW

Mass = 0.25 mol × 58.44 g/mol = 14.61 g

Therefore, the chemist should weigh 14.61 grams of NaCl.

Pedagogical Explanation:

This problem demonstrates the practical application of molecular weight in laboratory work. The relationship between moles, mass, and molecular weight is fundamental for preparing solutions and conducting experiments. This calculation bridges theoretical chemistry with practical laboratory work.

Key Definitions:

Mole: Amount of substance containing Avogadro's number of particles

Molecular Weight: Mass per mole of substance

Mass Calculation: Converting between moles and grams

Important Rules:

• Mass = moles × MW

• Moles = mass / MW

• Essential for laboratory preparation

Tips & Tricks:

• Use analytical balances for accurate weighing

• Consider purity of chemicals

• Store chemicals properly after weighing

Common Mistakes:

• Using wrong formula (MW/moles instead of moles×MW)

• Not accounting for hydrates

• Forgetting to use molecular weight

Question 4: Application-Based Problem - Empirical Formula

A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. What is its empirical formula? (Atomic masses: C=12.01, H=1.008, O=16.00)

Solution:

Step 1: Assume 100g sample and convert to moles

Carbon: 40.0g / 12.01 g/mol = 3.33 mol

Hydrogen: 6.7g / 1.008 g/mol = 6.65 mol

Oxygen: 53.3g / 16.00 g/mol = 3.33 mol

Step 2: Find the simplest whole number ratio

Divide by smallest number (3.33):

C: 3.33 / 3.33 = 1.0

H: 6.65 / 3.33 = 2.0

O: 3.33 / 3.33 = 1.0

Step 3: Write empirical formula

CH₂O is the empirical formula.

Therefore, the empirical formula is CH₂O.

Pedagogical Explanation:

This problem demonstrates how molecular weight calculations are used to determine chemical formulas from experimental data. The process involves converting mass percentages to moles and finding the simplest whole number ratio. This is fundamental in analytical chemistry.

Key Definitions:

Empirical Formula: Simplest whole number ratio of elements

Molecular Formula: Actual number of atoms in molecule

Composition Analysis: Determining elemental makeup

Important Rules:

• Convert mass % to moles

• Divide by smallest number to get ratios

• Multiply to get whole numbers if needed

Tips & Tricks:

• Assume 100g sample for percentage calculations

• Check that ratios are close to whole numbers

• Verify the calculation by checking mass percents

Common Mistakes:

• Not converting to moles first

• Forgetting to divide by smallest number

• Not checking if ratios are reasonable

Question 5: Multiple Choice - Isotope Effect

What is the molecular weight of water made with deuterium (²H) instead of regular hydrogen (¹H)? (Atomic masses: ²H=2.014, O=15.999)

Solution:

The answer is B) 20.027 g/mol. For D₂O (deuterium oxide): MW = (2 × 2.014) + (1 × 15.999) = 4.028 + 15.999 = 20.027 g/mol. This is heavier than regular water (H₂O) which has MW = 18.015 g/mol.

Pedagogical Explanation:

This question explores the effect of isotopes on molecular weight. Deuterium is an isotope of hydrogen with one neutron (compared to regular hydrogen which has no neutrons). Isotopic substitution is important in nuclear chemistry and tracer studies.

Key Definitions:

Isotope: Atoms of same element with different neutrons

Isotope: Atoms of same element with different neutrons

Deuterium: Heavy isotope of hydrogen (²H)

Isotopic Substitution: Replacing atoms with isotopes

Important Rules:

• Isotopes have same atomic number, different mass

• Molecular weight changes with isotope substitution

• Natural abundance affects average atomic mass

Tips & Tricks:

• Use exact isotope masses for precision

• Consider natural abundance for average masses

• Isotopes affect physical properties more than chemical

Common Mistakes:

• Using average atomic mass instead of isotope mass

• Forgetting to account for all atoms in formula

• Confusing atomic number with mass number

FAQ

Q: What's the difference between molecular weight and formula weight?

A: Molecular weight refers to the sum of atomic masses in a molecule (covalently bonded species), while formula weight refers to the sum of atomic masses in the formula unit of an ionic compound.

For example:

  • Water (H₂O) has a molecular weight of 18.015 g/mol
  • Sodium chloride (NaCl) has a formula weight of 58.44 g/mol

Both are calculated the same way: MW = Σ(atomic_mass × subscript), but the terminology reflects the nature of the bonding in the compound. The values are numerically identical, but the conceptual difference is important for understanding chemical structure.

Q: How do I handle complex formulas like CuSO₄·5H₂O?

A: For hydrated compounds like CuSO₄·5H₂O (copper sulfate pentahydrate), treat the water molecules as separate units:

\[ \text{MW} = \text{MW(CuSO₄)} + 5 \times \text{MW(H₂O)} \]

Calculate each component:

  • Cu: 1 × 63.55 = 63.55 g/mol
  • S: 1 × 32.07 = 32.07 g/mol
  • O: 4 × 16.00 = 64.00 g/mol
  • 5H₂O: 5 × 18.015 = 90.075 g/mol

Total MW = 63.55 + 32.07 + 64.00 + 90.075 = 249.695 g/mol

The dot (·) indicates water of hydration, not a covalent bond. Always account for all atoms in the formula unit.

About

Chemistry Team
This molecular weight calculator was created
This calculator was created by our General & Utility Calculators Team , may make errors. Consider checking important information. Updated: April 2026.