Molarity Calculator
Calculate molar concentration (molarity, M) of a solution from moles of solute and total solution volume in litres. The most common concentration unit in chemistry, essential for stoichiometry, titrations, dilutions, and any lab calculation involving solutions.
Last updated: May 2026
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About this calculator
Molarity (M) is defined as the number of moles of solute per litre of total solution: M = n/V, where n is moles (mol) and V is volume (L). The unit is mol/L, abbreviated M. The formula gives a precise, reproducible concentration that scales linearly: doubling the moles doubles the molarity; halving the volume doubles the molarity. Key conceptual points: it is per litre of SOLUTION (the final mixture), not per litre of solvent — adding 1 mole of NaCl to 900 mL of water gives slightly more than 1 L of solution, and the molarity is moles divided by that final volume. Molarity is temperature-dependent because the volume of liquid changes with temperature — water expands ~0.02% per °C near room temperature, so a 1.000 M solution at 20 °C is approximately 0.998 M at 30 °C. For temperature-independent concentration units use molality (mol/kg solvent) or mole fraction. Edge cases: very dilute solutions (μM, nM) use the same formula but require serial dilution to prepare; very concentrated solutions approach the solubility limit, where the assumption of complete dissolution breaks down. V = 0 gives undefined molarity (divide by zero). For mixed solvents (e.g., ethanol in water) the volume is still total solution volume but volume contraction becomes noticeable (mixing 50 mL ethanol + 50 mL water yields only ~96 mL). For ionic compounds that dissociate, molarity refers to the formula unit, not individual ions — 1 M NaCl is 1 M in formula units but 2 M in total ions (1 M Na⁺ + 1 M Cl⁻).
How to use
Example 1 — Preparing a salt solution. You weigh out 14.61 g of NaCl (MW 58.44 g/mol) and dissolve it in enough water to make 250 mL of solution. Convert to moles: n = 14.61/58.44 = 0.250 mol. Convert volume to litres: V = 0.250 L. Enter moles = 0.250 and volume = 0.250. M = 0.250/0.250 = 1.000 M. ✓ That is a standard 1 M saline stock — each litre contains 58.44 g of dissolved salt. Example 2 — A dilute biological buffer. You dissolve 0.0042 mol of HEPES (MW 238.3 g/mol) into 200 mL of pH 7.4 buffer. V = 0.200 L. Enter moles = 0.0042 and volume = 0.200. M = 0.0042/0.200 = 0.021 M = 21 mM. ✓ A typical biological-buffer concentration — strong enough to maintain pH against small acid/base additions but dilute enough not to interfere with most enzyme reactions or cell biology experiments.
Frequently asked questions
Why divide by solution volume rather than solvent volume?
Molarity is defined per litre of total solution because that is what you actually measure with a graduated cylinder, beaker, or volumetric flask — the total volume of the final mixed liquid. Using solvent volume instead would require knowing the partial volume of each component, which is impractical for everyday lab work. The distinction matters most for concentrated solutions: dissolving 100 g of salt in 1 L of water doesn’t give 1 L of solution — the dissolved salt adds volume, ending up around 1.04 L. If you computed molarity as ‘moles per litre of water added’ you would overstate the concentration by 4%. Using total solution volume keeps the math consistent regardless of concentration and matches the practical procedure of ‘dissolve solute, then dilute to the mark in a volumetric flask’.
What is the difference between molarity and molality?
Molarity (M) is moles of solute per litre of solution. Molality (m) is moles of solute per kilogram of solvent. The two are nearly identical for dilute aqueous solutions at room temperature — 1 L of water has a mass of ~1 kg, so 1 M ≈ 1 m for dilute aqueous solutions. The values diverge for concentrated solutions, non-aqueous solvents, or temperatures far from 25 °C. Crucially, molality is temperature-independent (mass doesn’t change with temperature) while molarity is temperature-dependent (volume does). For colligative properties — freezing-point depression, boiling-point elevation, osmotic pressure — molality is the correct unit because the proportionality constants are derived in mass terms. For routine bench work in mostly-aqueous solutions, molarity is what virtually everyone uses.
How do I prepare a specific molarity solution from a solid?
Three steps: (1) Compute moles needed: n = M × V, where M is target molarity and V is desired final volume in litres. For 500 mL of 0.1 M solution, n = 0.1 × 0.500 = 0.050 mol. (2) Convert moles to grams: g = n × MW, where MW is the molar mass. For glucose (MW 180.16): g = 0.050 × 180.16 = 9.008 g. (3) Weigh out that mass, add it to a 500 mL volumetric flask, fill with solvent to ¾ full, mix until dissolved, then top up exactly to the 500 mL line. The ‘top up to the line’ step is critical — adding the solute first changes total volume, so you cannot just dissolve in 500 mL of water and call it 500 mL of solution. Volumetric flasks are calibrated for total solution volume, which is exactly what molarity uses.
What are the most common mistakes people make with molarity?
The first is using millilitres instead of litres — a 1 mol / 100 mL calculation gives ‘0.01 M’ if you forget to convert, when it is actually 10 M. Always work in litres or remember 1 mL = 0.001 L. The second is preparing solutions by ‘dissolving in 1 L of water’ instead of ‘dissolving and diluting to 1 L total’ — slightly less concentrated than intended because the solute itself takes up volume. The third is forgetting that ionic compounds dissociate: 1 M MgCl₂ is 1 M in formula units but 3 M in dissolved ions (1 M Mg²⁺ + 2 M Cl⁻), which matters for ionic strength, conductivity, and osmotic calculations. The fourth is using volumetric flasks at temperatures different from their calibration (usually 20 °C); concentrated solutions change volume meaningfully with temperature. The fifth is confusing M (molarity, mol/L) with m (molality, mol/kg solvent) — they look similar but produce different numbers.
When should I not use this calculator?
Skip it for non-liquid solutions — gas mixtures use mole fractions or partial pressures, alloys and metallurgical samples use weight or atomic percent, and solid solutions like mineral phases have their own concentration conventions. Avoid it for solutions where the solute doesn’t fully dissolve (saturated or supersaturated mixtures) — effective molarity depends on what is actually in solution versus precipitated. It is the wrong unit for colligative-property calculations (freezing-point depression, boiling-point elevation, osmotic pressure) where you need molality. Do not use it for very dilute environmental samples in trace analysis where parts-per-million (ppm) or parts-per-billion (ppb) are more meaningful units. Finally, do not use it for biochemical work involving high-molecular-weight polymers or proteins where mass-per-volume (mg/mL or g/L) is the standard reporting unit, because exact molar masses can be hard to determine and the ‘mole’ concept loses precision for heterogeneous mixtures.