Mole Calculator

How to use this calculator

The first thing to do is choose a mode from the Calculate From dropdown. There are 3 options, and they correspond to 3 different starting situations.

Mode 1: Mass + Molar Mass (n = m / M)

Use this when you know the mass of a substance in grams and its molar mass in g/mol.

• Mass (m): Enter the mass of your sample in grams. If you have 18.015 g of water, enter 18.015.
• Molar Mass (M): Enter the molar mass of the substance in g/mol. For water (H₂O), that’s 18.015 g/mol. You can look this up from a periodic table by summing the atomic masses of all atoms in the formula.

This is the most common mode. Most lab problems give you a mass and ask for moles.

Mode 2: Number of Particles (n = N / Nₐ)

Use this when you know how many individual atoms or molecules you have.

• Number of Particles (N): Enter the count of atoms or molecules. The default shown is 6.02214076 × 10²³, which is exactly 1 mole. Enter your value in scientific notation or as a plain number.

Avogadro’s number (Nₐ = 6.02214076 × 10²³) is the denominator. The calculator handles that automatically.

Mode 3: Gas Volume at STP (n = V / 22.414)

Use this for ideal gases at standard temperature and pressure (0°C, 1 atm).

• Gas Volume (V): Enter the volume of gas in liters. At STP, 1 mole of any ideal gas occupies 22.414 L. Enter your measured volume and the calculator divides by that molar volume.

This mode only applies to gases at STP. Room temperature (25°C) uses a different molar volume (24.465 L/mol), so don’t use Mode 3 for gases measured at room conditions.

Click Calculate to get your moles. Click Clear to reset all fields.

What problem this calculator solves

Every stoichiometry problem in chemistry eventually needs moles. Reaction equations are written in molar ratios: 2H₂ + O₂ → 2H₂O means 2 moles of hydrogen react with 1 mole of oxygen. To use that ratio, you need to convert your actual quantities (grams, liters, particle counts) into moles first.

Without moles, the ratios don’t apply. You can’t say “2 grams of hydrogen reacts with 1 gram of oxygen” because the atomic masses are different. But “2 moles of hydrogen” and “1 mole of oxygen” is exact.

The other place this shows up constantly is in concentration problems. Molarity (mol/L) requires moles. Mole fraction requires moles. Virtually every quantitative chemistry concept routes through this unit at some point.

The calculator makes the entry step fast so you can focus on what comes after.

The concept explained simply

Avogadro figured out that equal volumes of gases at the same temperature and pressure contain the same number of molecules, regardless of what the gas is. That’s a remarkable fact. It means the “amount” of a substance can be described independently of what the substance is.

That amount is the mole. One mole of anything contains 6.02214076 × 10²³ of that thing. One mole of carbon atoms. One mole of water molecules. One mole of electrons. The count is the same.

What changes is the mass. Carbon atoms are heavier than hydrogen atoms, so 1 mole of carbon (12.011 g) weighs more than 1 mole of hydrogen gas (2.016 g). The molar mass is just the gram equivalent of that count. It’s the bridge between “how many” and “how much.”

One mole is to atoms what a dozen is to eggs. The only difference is scale: 6.02 × 10²³ instead of 12, and the "unit" on the scale is the atomic mass unit instead of nothing.

The three formulas explained

Each calculator mode uses a different formula depending on what you’re starting with.

Mode 1: from mass

n is moles, m is mass in grams, M is molar mass in g/mol. Divide grams by grams-per-mole and you get moles. The units cancel cleanly: (g) / (g/mol) = mol.

Mode 2: from particle count

N is the number of particles (atoms or molecules), Nₐ is Avogadro’s number (6.02214076 × 10²³ mol⁻¹). Divide your particle count by Avogadro’s number to get moles.

Mode 3: from gas volume at STP

V is the volume of gas in liters, 22.414 L/mol is the molar volume of an ideal gas at STP (0°C, 1 atm). Divide volume by 22.414 to get moles.

All 3 formulas are solving for the same thing. They just start from different knowns. The mole is the output in every case.

Reference table: molar masses for common substances

You need the molar mass for Mode 1. Here are values for substances that come up constantly in general chemistry.

Substance	Formula	Molar mass (g/mol)
Water	H₂O	18.015
Sodium chloride	NaCl	58.443
Carbon dioxide	CO₂	44.010
Glucose	C₆H₁₂O₆	180.156
Sulfuric acid	H₂SO₄	98.079
Ammonia	NH₃	17.031
Oxygen gas	O₂	31.999
Hydrogen gas	H₂	2.016
Methane	CH₄	16.043
Calcium carbonate	CaCO₃	100.087
Iron	Fe	55.845
Ethanol	C₂H₅OH	46.068

For anything not listed, sum the atomic masses from the periodic table. Carbon = 12.011, Hydrogen = 1.008, Oxygen = 15.999, Nitrogen = 14.007, and so on.

Real-world examples

The lab preparation problem

A student needs to prepare a 0.5 mol/L solution of NaCl and wants to make 250 mL. How many grams of NaCl do they need?

The gas cylinder scenario

A technician has a 5.0 L sample of nitrogen gas collected at STP. How many moles is that, and how many molecules?

The reaction stoichiometry check

A combustion reaction burns 44 g of propane (C₃H₈, molar mass = 44.097 g/mol). The balanced equation is C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. How many moles of CO₂ are produced?

Common mistakes people make

Mixing up mass and molar mass. Mass (m) is specific to your sample. Molar mass (M) is a property of the substance itself. If you have 5 g of NaCl, the mass is 5 g and the molar mass is 58.443 g/mol, always, regardless of how much you have. Swapping these gives you a nonsense answer.

Using Mode 3 at the wrong temperature. This comes up constantly. A student measures 2.5 L of oxygen at room temperature (25°C), plugs it into Mode 3, and gets n = 2.5 / 22.414 = 0.111 mol. The correct answer using 24.465 L/mol is 0.102 mol. For a quick estimate that’s fine. For a graded lab report, it’s a mistake. Check your conditions before choosing Mode 3.

Counting formula units vs. atoms. Mode 2 asks for atoms or molecules depending on what you’re tracking. 1 mole of NaCl gives you 6.022 × 10²³ formula units, but 1.204 × 10²⁴ ions (Na⁺ and Cl⁻ combined). If your problem asks for moles of ions, not moles of NaCl, the answer differs by a factor of 2. The calculator gives you moles of whatever particle count you enter.

Forgetting to convert units before entering. The mass field expects grams. If your mass is in milligrams, divide by 1000 first. If it’s in kilograms, multiply by 1000. The molar mass field expects g/mol. These aren’t flexible.

Computing molar mass incorrectly. CaCO₃ has 1 calcium (40.078), 1 carbon (12.011), and 3 oxygens (3 × 15.999 = 47.997). Total: 100.086 g/mol. A common error is counting only 1 oxygen. Double-check subscripts in the formula.

Hidden factors most people ignore

Hydrated salts have extra mass. CuSO₄·5H₂O (blue vitriol) has a molar mass of 249.68 g/mol, not 159.61 g/mol (anhydrous CuSO₄). If you weigh out a hydrated salt and use the anhydrous molar mass, you’ll undercount your moles. Always check whether your compound is hydrated.

Isotopic composition matters at high precision. The molar mass of carbon is 12.011 g/mol, not exactly 12, because about 1.1% of natural carbon is carbon-13. For most chemistry, this is negligible. For isotope-specific work or mass spectrometry, the distinction matters.

Purity affects actual moles. A bottle labeled “58.44 g of NaCl” might be 99.2% pure. That means you actually have 57.96 g of NaCl and 0.47 g of something else. In a teaching lab this rarely matters. In analytical work, it does. Reagent grade vs. lab grade vs. technical grade have different purity levels.

Avogadro’s number was redefined in 2019. The SI system redefined the mole so that Avogadro’s constant is exactly 6.02214076 × 10²³ mol⁻¹, with no uncertainty. Before that, it was a measured value. For all practical purposes this changes nothing, but it’s worth knowing if you see a slightly different value in older textbooks.

The mole isn't arbitrary. It was chosen so that the molar mass in g/mol numerically equals the atomic mass in atomic mass units. That alignment is what makes the m/M conversion exact.

What to do with the result

For stoichiometry: once you have moles of your starting material, multiply by the molar ratio from the balanced equation. If the equation says 1:2, you multiply by 2. If it says 3:1, you multiply by 1/3. Then convert back to grams if needed using n × M.

For solution preparation: if you know the molarity (C) and volume (V) you want, the moles you need is n = C × V. Use Mode 1 in reverse: grams needed = n × M. Weigh that mass, dissolve, dilute to the target volume.

For particle-level problems: multiply your moles by 6.022 × 10²³ to get the number of atoms or molecules. This comes up in statistical mechanics and when problems ask “how many molecules are in…”

For ideal gas problems: moles connects to pressure and temperature via PV = nRT. Once you have n from the calculator, you can find pressure, volume, or temperature given the other two.

For mole fraction: this is where the mole calculator and the mole fraction calculator work together. Use this calculator to convert each component from grams to moles. Then enter those mole values into the mole fraction calculator.

Limitations and misconceptions

The calculator assumes pure substances. If your sample is a mixture, you can’t use a single molar mass for the whole thing. You’d need to know the composition, separate out the components, and calculate moles for each individually.

Mode 3 assumes ideal gas behavior. Real gases deviate from ideal behavior at high pressures and low temperatures. For most intro chemistry problems at or near room conditions, this deviation is small enough to ignore. For high-pressure industrial applications, it’s not.

A persistent misconception is that the mole is somehow a physical object or a container. It’s a count. “One mole of water” is just a name for 6.022 × 10²³ water molecules. Nothing special happens at that exact quantity. The number was chosen because it lines up the atomic mass scale with the gram scale.

The calculator also doesn’t help you find molar mass. If you don’t know the molar mass of your compound, you’ll need a periodic table and the molecular formula. The molar mass reference table in this article covers the most common cases.

The bottom line

The mole is the unit that makes chemistry quantitative. Without it, balanced equations are just symbolic. With it, you can predict how many grams of product a reaction will yield, how much reagent to weigh out for a solution, or how many molecules are in a gas sample.

This calculator handles the 3 most common entry points: mass, particle count, and gas volume at STP. Pick the mode that matches your starting information, enter the numbers, and you have moles in one step.

Get the moles right and everything downstream works. Get them wrong and every calculation after this point is off by the same factor.