Molarity Calculator
Calculate the molarity of a solution from mass, molar mass, and volume — or find the mass of solute needed to make a solution of a target concentration.
How to use this calculator
- 1Choose whether you know the mass and want the molarity, or know the target molarity and want the mass.
- 2Enter the molar mass of the solute (the sum of its atomic masses).
- 3Enter the solution volume in litres or millilitres.
- 4Read the molarity or the mass to weigh out, along with the moles involved.
How it works
Molarity
molarity (M) = moles of solute ÷ litres of solution moles = mass ÷ molar mass so M = mass ÷ (molar mass × volume in L) mass needed = molarity × volume × molar mass
Molarity is the most common way chemists express concentration: the number of moles of a dissolved substance per litre of solution. Because a mole is a fixed number of particles, molarity tells you how many solute particles are present in a given volume, which is exactly what reactions depend on. To find it, convert the mass of solute to moles by dividing by its molar mass (the sum of the atomic masses in its formula), then divide by the volume in litres. The relationship rearranges easily: to prepare a solution of a chosen molarity, multiply the target concentration by the volume and the molar mass to find the mass to weigh out. The one subtlety is that the volume is that of the finished solution, not the water added, because dissolving a solute changes the total volume.
Worked example
To make one litre of 1 M sodium chloride solution, you need 1 mol × 58.44 g/mol = 58.44 g of NaCl, dissolved and made up to 1 litre. Conversely, dissolving 58.44 g of NaCl and topping up to 1 litre gives a 1 mol/L solution — the concentration of roughly normal saline is about a sixth of this.
Molarity Calculator: the complete guide
What a mole and molarity really mean
Chemistry cares about the number of particles that react, not their mass, because reactions happen atom-to-atom and molecule-to-molecule. But counting individual particles is impossible, so chemists use the mole — a fixed, enormous count (about 6.022 × 10²³) that bridges the invisible world of atoms and the measurable world of grams. The molar mass of a substance, in grams per mole, is numerically the same as its formula weight in atomic mass units, which is what makes the conversion between mass and moles a simple division.
Molarity then packages this into a concentration: moles per litre. When a recipe or reaction calls for a 0.1 M solution, it is specifying a particle count per unit volume, so that a measured volume delivers a known number of reacting particles. This is why molarity, rather than a mass-based concentration like grams per litre, is the natural unit for stoichiometry — the volume of one solution reacting with another translates directly into a ratio of moles.
Preparing a solution correctly
A subtle but important point trips up many beginners: molarity is defined per litre of solution, not per litre of solvent. When a solute dissolves, it takes up space and changes the total volume, sometimes noticeably. So the correct way to make a precise 1 M solution is not to add the solute to a full litre of water — that would give slightly more than a litre and therefore a concentration a little below 1 M. Instead, dissolve the solute in a smaller amount of water, then add water up to the final volume mark.
This is exactly what a volumetric flask is for: it has a single calibrated line marking a precise volume. The solute is weighed accurately, dissolved in part of the solvent, and then topped up to the line. For everyday, non-critical work the difference is often negligible, but for analytical chemistry, where concentrations must be exact, the dissolve-then-dilute-to-volume method is essential. The calculator gives the mass to weigh; the technique is what turns that mass into an accurate concentration.
Molarity, molality, and dilution
Molarity is not the only concentration unit, and knowing where it falls short matters. Because it is defined per litre of solution, and volume changes with temperature, molarity drifts slightly as a solution warms or cools. For work where temperature varies, chemists sometimes prefer molality — moles per kilogram of solvent — which is based on mass and therefore temperature-independent. For most laboratory and educational purposes at room temperature, though, molarity is the standard.
Molarity also underpins dilution, one of the most frequent lab operations. The dilution equation, M₁V₁ = M₂V₂, says that the moles of solute are conserved when you add solvent, so a concentrated stock can be diluted to a working concentration by a simple proportion. Preparing a 0.1 M solution from a 1 M stock, for instance, means taking one part stock to nine parts water. Mastering molarity is thus the foundation not just for making solutions from solids, but for the routine dilutions that fill a working chemist's day.
Frequently asked questions
How do I calculate molarity?
Divide the moles of solute by the volume of solution in litres. To get moles, divide the solute's mass by its molar mass. Combined, molarity = mass ÷ (molar mass × litres). For example, 58.44 g of NaCl (molar mass 58.44) in 1 litre gives 1 mol/L.
How much solute do I need for a given molarity?
Multiply the target molarity by the volume in litres to get moles, then multiply by the molar mass to get grams. For 1 litre of 1 M NaCl: 1 mol/L × 1 L × 58.44 g/mol = 58.44 g. Dissolve it and make the solution up to the final volume.
Is molarity per litre of water or solution?
Per litre of solution, not solvent. Dissolving a solute changes the total volume, so a precise solution is made by dissolving the solute in some water and then topping up to the final volume mark, ideally in a volumetric flask — not by adding solute to a full litre of water.
What is the difference between molarity and molality?
Molarity is moles per litre of solution; molality is moles per kilogram of solvent. Molarity is more common but changes slightly with temperature because volume does. Molality, being mass-based, is temperature-independent and preferred when temperature varies significantly.