Grams to Moles Calculator
Convert mass in grams to moles using the molar mass of a substance. This calculator also shows the number of molecules or formula units using Avogadro's number (6.022 × 10²³). Select a common substance or enter a custom molar mass. See also our Moles to Atoms Calculator and Molar Mass Calculator for related computations.
How to Convert Grams to Moles
Converting between grams and moles is one of the most fundamental calculations in chemistry. The mole is the SI unit for amount of substance, and it provides the bridge between the macroscopic world (grams that you can weigh) and the microscopic world (individual atoms and molecules). Every stoichiometric calculation in chemistry requires this conversion at some point.
- Identify the substance and determine its molar mass (molecular weight) in g/mol.
- The molar mass equals the sum of atomic masses of all atoms in the formula.
- Measure or note the mass of the substance in grams.
- Divide the mass by the molar mass: moles = mass (g) / molar mass (g/mol).
- To find the number of molecules, multiply moles by Avogadro's number (6.022 × 10²³).
The molar mass is numerically equal to the atomic or molecular weight expressed in atomic mass units (amu or u), but with units of g/mol. For example, water (H₂O) has a molecular weight of 18.015 amu, so its molar mass is 18.015 g/mol. This means that 18.015 grams of water contains exactly one mole (6.022 × 10²³) of water molecules.
Grams to Moles Formula
Moles = Mass (g) / Molar Mass (g/mol)
n = m / M
Number of molecules = Moles × Avogadro's Number
N = n × Nₐ
N = n × 6.02214076 × 10²³
Reverse (moles to grams):
Mass (g) = Moles × Molar Mass (g/mol)
Where:
n = number of moles (mol)
m = mass of substance (g)
M = molar mass (g/mol)
N = number of particles (molecules, atoms, or formula units)
Nₐ = Avogadro's number = 6.02214076 × 10²³ mol⁻¹
The formula is a simple division, but its significance is profound. It connects the measurable quantity (mass) to the chemically meaningful quantity (moles, which represent a specific number of particles). One mole of any substance contains exactly 6.02214076 × 10²³ particles — this is the defined value of Avogadro's number as of the 2019 SI redefinition.
Example Calculation
Problem: How many moles and molecules are in 18 g of water (H₂O)?
Given:
• Mass = 18 g
• Molar mass of H₂O = 2(1.008) + 15.999 = 18.015 g/mol
Solution:
Moles = 18 g / 18.015 g/mol = 0.9992 mol ≈ 1.00 mol
Molecules = 0.9992 × 6.022 × 10²³ = 6.017 × 10²³ molecules
Answer: 18 g of water is approximately 1 mole, containing about 6.02 × 10²³ water molecules.
Additional Example: How many moles in 100 g of NaCl (MW = 58.44)?
• Moles = 100 / 58.44 = 1.711 mol
• Formula units = 1.711 × 6.022 × 10²³ = 1.030 × 10²⁴
Molar Mass Reference Table
| Substance | Formula | Molar Mass (g/mol) | 1 mol = ? grams |
|---|---|---|---|
| Water | H₂O | 18.015 | 18.015 |
| Sodium chloride | NaCl | 58.44 | 58.44 |
| Glucose | C₆H₁₂O₆ | 180.16 | 180.16 |
| Ethanol | C₂H₅OH | 46.07 | 46.07 |
| Carbon dioxide | CO₂ | 44.01 | 44.01 |
| Sulfuric acid | H₂SO₄ | 98.079 | 98.079 |
| Sodium hydroxide | NaOH | 40.00 | 40.00 |
| Calcium carbonate | CaCO₃ | 100.09 | 100.09 |
| Ammonia | NH₃ | 17.031 | 17.031 |
| Oxygen gas | O₂ | 31.998 | 31.998 |
| Nitrogen gas | N₂ | 28.014 | 28.014 |
| Methane | CH₄ | 16.043 | 16.043 |
Frequently Asked Questions
What is a mole in chemistry?
A mole is the SI unit for amount of substance. One mole contains exactly 6.02214076 × 10²³ elementary entities (atoms, molecules, ions, or other particles). It is the chemist's "counting unit" — just as a dozen means 12, a mole means 6.022 × 10²³. The mole bridges the gap between individual atoms (too small to count) and laboratory quantities (grams that can be weighed).
What is Avogadro's number?
Avogadro's number (Nₐ) is 6.02214076 × 10²³ mol⁻¹. It is the number of particles in one mole of any substance. Named after Amedeo Avogadro, this constant was redefined in 2019 as an exact value (no uncertainty). It connects the atomic mass unit (amu) to the gram: 1 amu = 1 g / Nₐ = 1.66054 × 10⁻²⁴ g.
How do I find the molar mass of a compound?
To find the molar mass, add up the atomic masses of all atoms in the chemical formula. For H₂O: 2 × H (1.008) + 1 × O (15.999) = 18.015 g/mol. For Ca(OH)₂: 1 × Ca (40.078) + 2 × O (15.999) + 2 × H (1.008) = 74.092 g/mol. Atomic masses are found on the periodic table. Use our Molar Mass Calculator for complex formulas.
What is the difference between molecular weight and molar mass?
Molecular weight (or molecular mass) is the mass of one molecule expressed in atomic mass units (amu or Da). Molar mass is the mass of one mole of molecules expressed in grams per mole (g/mol). They are numerically equal but have different units. For example, water has a molecular weight of 18.015 amu and a molar mass of 18.015 g/mol.
How many molecules are in 1 gram of water?
Moles of water = 1 g / 18.015 g/mol = 0.05551 mol. Number of molecules = 0.05551 × 6.022 × 10²³ = 3.343 × 10²² molecules. This means even a single gram of water contains over 33 sextillion molecules — an incomprehensibly large number that illustrates why chemists use moles rather than counting individual molecules.
Can I convert grams to moles for elements?
Yes, the same formula applies. For elements, the molar mass equals the atomic mass from the periodic table. For example, 55.85 g of iron (Fe, atomic mass = 55.845) equals 1 mole of iron atoms. For diatomic elements (H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂), use the diatomic molar mass (e.g., O₂ = 31.998 g/mol, not 15.999).
The Mole Concept in Chemistry
The mole is arguably the most important concept in chemistry. It provides the quantitative foundation for all chemical calculations, from simple stoichiometry to complex thermodynamic analyses. Without the mole concept, it would be impossible to predict how much product a reaction will yield, how much reagent to use, or how concentrated a solution is.
The historical development of the mole concept parallels the development of atomic theory. In the early 19th century, John Dalton proposed that elements consist of atoms with characteristic masses. Avogadro hypothesized that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. These ideas eventually led to the mole as a practical unit for counting atoms and molecules by weighing macroscopic samples.
In laboratory practice, the grams-to-moles conversion is performed dozens of times daily. When preparing solutions, chemists weigh a substance in grams and convert to moles to determine concentration. When running reactions, they convert reactant masses to moles to determine stoichiometric ratios and identify limiting reagents. When analyzing products, they convert measured masses to moles to calculate yields and purities.
The 2019 redefinition of the SI system fixed Avogadro's number at exactly 6.02214076 × 10²³, making the mole independent of the kilogram. Previously, the mole was defined as the number of atoms in exactly 12 grams of carbon-12. The new definition is conceptually simpler and more precise, though it has no practical impact on everyday chemistry calculations.
Understanding the scale of Avogadro's number helps appreciate why the mole is necessary. If you could count atoms at a rate of one million per second, it would take about 19 billion years (longer than the age of the universe) to count one mole. This enormous number is what makes individual atoms invisible to our senses while allowing mole-scale quantities to be easily handled in the laboratory.