Calculate molecular weight from chemical formula. Supports elements, compounds, hydrates, and polymers with a full periodic table database.
The molecular weight calculator determines the molar mass of any chemical compound from its elemental composition. By summing the atomic masses of all atoms in a molecular formula, you obtain the molecular weight in grams per mole (g/mol) — the mass of one mole or 6.022 × 10²³ molecules of the substance.
Molecular weight is arguably the most frequently needed value in chemistry. It appears in virtually every calculation: converting grams to moles, determining molarity, calculating yields, performing dilutions, and balancing equations quantitatively. Having an accurate molecular weight is the first step in almost any stoichiometry problem.
This calculator accepts common chemical formulas, breaks them down into their constituent elements, and computes the total molecular weight using IUPAC-recommended atomic masses. It also shows the percent composition of each element — useful for empirical formula determination and analytical chemistry.
For best results, combine calculator output with direct observation and periodic check-ins with a veterinarian or qualified advisor. Small adjustments made early usually improve comfort, safety, and long-term outcomes more than large corrective changes made later.
This calculator provides instant molecular weight and percent composition calculations for any compound. It eliminates the need to look up individual atomic masses and perform tedious summations. This molecular weight calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
Molecular Weight = Σ(number of atoms × atomic mass) for all elements For example, H₂O = 2(1.008) + 1(16.00) = 18.015 g/mol Percent Composition = (atoms × atomic mass / MW) × 100%
Result: 98.079 g/mol
For H₂SO₄: H = 2(1.008) = 2.016, S = 1(32.06) = 32.06, O = 4(16.00) = 64.00. Total = 2.016 + 32.06 + 64.00 = 98.079 g/mol.
Standard atomic weights are published by IUPAC and updated periodically as measurements improve. These weighted averages account for the natural isotopic distribution of each element. For example, chlorine's atomic weight is 35.45, not 35 or 37, because it's about 75% ³⁵Cl and 25% ³⁷Cl in nature.
If you know the percent composition from elemental analysis, you can work backward to find the empirical formula. Divide each element's mass percentage by its atomic weight to get mole ratios, then reduce to the smallest whole numbers. This is the reverse of the molecular weight calculation.
Mass spectrometry measures molecular weight with extreme precision. The molecular ion peak in a mass spectrum corresponds to the monoisotopic mass (using the most abundant isotope of each element), which differs slightly from the standard molecular weight. High-resolution mass spectrometry can determine molecular formulas by matching exact masses to within fractions of a Dalton.
Molecular weight is dimensionless (relative to ¹²C), while molar mass has units of g/mol. Numerically they are identical. In practice, chemists use the terms interchangeably.
The same way as for molecular compounds: sum the atomic weights of all atoms in the formula unit. NaCl = 22.99 + 35.45 = 58.44 g/mol. For ionic compounds, we more accurately call it "formula weight."
Include the water of crystallization. For CuSO₄·5H₂O: CuSO₄ = 159.61 + 5 × 18.015 = 249.69 g/mol. Always specify whether your compound is hydrated or anhydrous.
Use the standard atomic weights from IUPAC, which represent the weighted average of all naturally occurring isotopes. These are the numbers on the periodic table, like 12.011 for carbon.
For most calculations, 2-4 decimal places suffice. For high-precision work like mass spectrometry or exact mass determination, use monoisotopic masses (mass of the most abundant isotope).
Percent composition shows the mass percentage of each element, essential for determining empirical formulas from experimental data, checking compound purity, and analytical chemistry. This keeps planning practical and lowers the chance of preventable errors.