Molecular weight from any formula — type H2O and get 18.015 g/mol with the breakdown.
Every stoichiometry problem in chemistry runs through a single tollbooth: the molar mass. Your balance measures grams, but chemical equations think in moles, and the molar mass — the mass of one mole of a substance, in grams per mole — is the conversion factor between those two worlds. Without it you cannot turn “25.0 grams of sodium chloride” into the number of formula units that will actually react, and every limiting-reactant, percent-yield, and molarity calculation stalls at the first step.
Computing a molar mass by hand is simple but tedious: look up each element’s atomic mass on the periodic table, multiply by the subscript in the formula, and add everything together. For water, H2O, that means two hydrogens at 1.008 plus one oxygen at 15.999, giving 18.015 g/mol. For something like calcium carbonate, CaCO3, the arithmetic grows — and with long formulas or hydrates, transcription slips creep in. That is the entire job this calculator does: parse any formula you type, multiply each element’s count by its atomic mass, and show the full breakdown so you can verify every line.
Type a formula such as H2O, NaCl, or CaCO3 into the box and press Calculate Molar Mass. The result panel lists each element’s contribution — for example “H × 2 = 2.016” — then totals them to three decimal places. The sections below explain what molar mass really means, work through examples by hand, and cover the traps (diatomic elements, waters of hydration, isotopic averages) that catch students every semester.
Chemists count particles in moles because atoms are far too small to count individually: one mole contains 6.022 × 1023 entities, Avogadro’s number. But no laboratory instrument counts particles; the balance weighs grams. Molar mass is the exchange rate between these two currencies. It answers the question “how heavy is a mole of this stuff?” — and because a mole of carbon-12 is defined to weigh exactly 12 grams, every other molar mass is measured relative to that anchor.
The atomic masses on the periodic table are not the mass of a single atom but weighted averages over each element’s naturally occurring isotopes. Chlorine is listed at 35.45, not 35 or 37, because natural chlorine is roughly three parts chlorine-35 to one part chlorine-37. This is why molar masses usually have decimals, and it is also why the molar mass of a compound made from elements with very mixed isotopes can vary slightly by source — the differences are tiny, but they exist, and metrologists who need extreme precision use isotopically characterized materials.
A crucial distinction: molar mass (g/mol) is a property of a mole of substance, while molecular weight is technically dimensionless — the mass of one molecule relative to 1/12 of a carbon-12 atom. In practice, working chemists use the two terms interchangeably and report both as a number with units g/mol. The calculator follows that working convention: it computes the relative molecular mass from standard atomic weights and presents it as a molar mass in g/mol, which is what every stoichiometry calculation needs.
Finally, molar mass is what makes formulas comparable. Saying “I used 10 grams of reactant A and 10 grams of reactant B” tells you almost nothing about the chemistry, because 10 grams of hydrogen is five moles while 10 grams of lead is a twentieth of a mole. Converting both to moles via molar mass reveals the true reacting ratio, which is why every serious quantitative step in chemistry — limiting reactants, titrations, solution preparation — begins with this one number.
Key terms:
Here n is the subscript of each element in the formula and A is its standard atomic mass. Add one term per element and the sum is the molar mass. Once you have it, the second line converts any weighed mass into moles — the direction of travel for nearly all of stoichiometry. Going the other way, multiply moles by molar mass to predict the mass a reaction should produce, which is exactly the theoretical-yield step in a percent yield calculation.
Table salt is the classic first example: one sodium, one chlorine, no subscripts to trip over.
Because NaCl is ionic, chemists say “formula mass” rather than “molecular mass,” but the number and its use are identical. The calculator treats every formula the same way.
Limestone, chalk, and antacid tablets: CaCO3 has three elements and a subscript that applies to oxygen only.
The classic student error here is multiplying the carbon by 3 along with the oxygen. Subscripts attach only to the element immediately before them — the 3 in CO3 counts oxygens, not carbons.
A lab technician must prepare 500 mL of 0.200 M sodium chloride solution. Molarity means moles per liter, so the needed moles are 0.200 mol/L × 0.500 L = 0.100 mol. Multiplying by the molar mass, 0.100 mol × 58.44 g/mol = 5.844 g. The technician weighs 5.84 g of NaCl, dissolves it, and dilutes to exactly 500 mL. Every solution ever prepared in every lab on earth runs through this exact chain, and the molar mass is the link that turns a concentration on paper into a mass on the balance.
Second scenario: a student titrates an unknown acid and finds 0.0250 mol of NaOH neutralized it in a 1:1 reaction. The unknown therefore contained 0.0250 mol of acid. If the acid turns out to be acetic acid, HC2H3O2, with molar mass 60.052 g/mol, the sample held 0.0250 × 60.052 = 1.501 g. Without the molar mass, the titration gives moles but the question “how many grams were in the sample” stays unanswered.
Third: an industrial buyer compares two suppliers of magnesium hydroxide, Mg(OH)2, for neutralizing acidic wastewater. The relevant number is neutralizing capacity per kilogram, which starts from the molar mass: Mg (24.305) + 2 × O (2 × 15.999) + 2 × H (2 × 1.008) = 58.319 g/mol, with each mole neutralizing two moles of acid. The buyer who skips this calculation and compares price per kilogram alone may buy the cheaper powder and get less neutralization per dollar.
Fourth: a forensic chemist identifies a white powder as glucose, C6H12O6, molar mass 180.156 g/mol, versus sucrose, C12H22O11, at 342.297 g/mol. A 1.00 g sample is 0.00555 mol of glucose but only 0.00292 mol of sucrose — nearly a factor of two difference in particle count from the same mass. In quantitative analysis, identifying the compound (and hence its molar mass) must come before any mass-to-moles conversion means anything.
Pharmaceutical manufacturing: Every batch record for a drug starts with molar masses: converting the prescribed moles of each starting material into kilograms to charge the reactor. A misplaced decimal in a molar mass propagates into every downstream quantity, which is why pharma calculations are independently double-checked — the arithmetic is simple, but the consequences of getting it wrong are not.
Environmental monitoring: Air-quality limits are set in micrograms per cubic meter, but gas analyzers often report parts per million by volume. Converting between them requires the pollutant’s molar mass (via the ideal gas law), so the molar mass of NO2 (46.006 g/mol) versus SO2 (64.066 g/mol) directly changes whether a reading exceeds the legal limit.
Food science: Nutrition labels, preservative limits, and acidity regulators are all specified by mass, but their chemical behavior depends on moles. Formulating a buffer for a beverage at a target pH means converting the desired molar concentrations of citric acid and its salts into grams per batch — molar mass arithmetic at production scale.
Metallurgy: Smelting calculations convert tons of ore into tons of metal using the molar masses of the oxide and the metal. The theoretical maximum metal from a shipment of hematite, Fe2O3 (159.687 g/mol), versus the iron it contains (2 × 55.845 = 111.69 g per mole of ore) sets the benchmark every furnace is judged against.
Forensics: Drug-quantity laws often hinge on mass thresholds, but pharmacology hinges on moles. Converting a seized mass into moles via molar mass lets toxicologists compare the potency of different substances on a per-molecule basis, which is the only comparison that reflects how the body actually encounters the drug.
Space exploration: Life-support and propulsion calculations are mass budgets, and every gram launched costs money. Knowing that electrolyzing water yields hydrogen at 2.016 g/mol lets mission planners compute exactly how much water must be carried or recycled to supply a crew’s oxygen for a given mission duration.
One pattern worth internalizing: molar mass errors are silent multipliers. Unlike a conceptual mistake, which usually produces an obviously absurd answer, a wrong molar mass gives a plausible-looking number that is wrong by a fixed ratio — and that ratio infects every downstream step: moles, limiting reactant, theoretical yield, percent yield. When a whole calculation chain disagrees with experiment by a clean factor, check the molar masses first.
History: the long road to reliable atomic weights. John Dalton published the first table of relative atomic weights in the early 1800s, but chemists disagreed for decades about basic formulas — was water HO or H2O? Without agreement, every molar mass was uncertain. The breakthrough came at the Karlsruhe Congress of 1860, where Stanislao Cannizzaro convinced the assembled chemists to adopt consistent atomic weights based on Avogadro’s hypothesis. Within a decade Mendeleev had his periodic table, and molar masses finally became the reliable conversion factors we use today. Every time this calculator totals a formula, it leans on that 1860 consensus.
A worked scenario with numbers: hydrate versus anhydrous. A student needs 0.0500 mol of copper(II) sulfate for an experiment and grabs the blue crystals — the pentahydrate, CuSO4 with 5 H2O. The anhydrous molar mass is 63.546 + 32.06 + 4 × 15.999 = 159.602 g/mol, but the hydrate adds 5 × 18.015 = 90.075, totaling 249.677 g/mol. Weighing 0.0500 × 159.602 = 7.98 g of the blue crystals delivers only 0.0320 mol — a 36 percent shortfall that ruins the experiment. The correct mass is 0.0500 × 249.677 = 12.48 g. Hydrates are the single most expensive molar-mass trap in teaching labs.
Edge cases and pro tips: isotopes and precision. Standard atomic weights are averages over natural isotopic abundances, which vary slightly by source — lithium from different mines, for instance, has measurably different isotopic ratios. For ordinary work this is negligible, but isotope geochemists and metrologists use source-specific masses. A related edge case: this calculator covers the twelve most common elements in teaching chemistry; formulas containing others (say, AgNO3) will report an unknown-element message rather than a wrong number, which is deliberate — a refused calculation beats a confident wrong one.
When molar mass is not the number you need. For mixtures — air, gasoline, alloys — there is no single molar mass, only an average weighted by composition, and using one anyway is a category error. For polymers, the “molar mass” is a distribution (number-average versus weight-average), not one number, because chains differ in length. And in mass spectrometry, what matters is the monoisotopic mass (lightest isotopes only), which differs slightly from the average-mass molar mass used in bench stoichiometry. Match the mass definition to the problem.
Molar mass = sum over each element of (subscript × standard atomic mass), in g/mol; e.g., H2O = 2(1.008) + 15.999 = 18.015 g/mol. It converts mass to moles (moles = grams ÷ molar mass) and is required for limiting-reactant, theoretical-yield, and solution-concentration calculations. Atomic masses are weighted averages over natural isotopes. Source: Sigma Calculator, Molar Mass Calculator, https://sigmacalculator.net/chemistry-calculators/molar-mass-calculator/
Cite this page as: Molar Mass Calculator, Sigma Calculator, https://sigmacalculator.net/chemistry-calculators/molar-mass-calculator/
Last updated: October 12, 2026