Molarity and Solution Concentration Explained

Molarity is the most widely used way to express the concentration of a solution in chemistry. Whether you are titrating an acid in a general chemistry lab, preparing reagents for a biochemistry experiment, or scaling up a pharmaceutical formulation, understanding molarity is essential. It connects the amount of dissolved substance to the volume of solution in a way that makes stoichiometric calculations straightforward and practical.

This guide explains what molarity means, how to calculate it from scratch, how to perform dilution calculations, and how to prepare solutions accurately in the laboratory.

What Is Molarity?

Molarity is a measure of solution concentration defined as the number of moles of solute dissolved in one liter of solution. It is denoted by a capital M and expressed in units of mol/L (moles per liter).

A 1 M (one molar) solution of sodium chloride contains exactly one mole of NaCl (58.44 grams) dissolved in enough water to make one liter of total solution. Notice the critical distinction: molarity is based on the total solution volume, not the volume of solvent added. This means you cannot simply add one mole of solute to one liter of water, because the dissolved solute will increase the total volume slightly beyond one liter.

The Molarity Formula

Since the number of moles is calculated from mass divided by molar mass, you can expand the formula:

This expanded form is the most practical version because laboratory measurements typically involve weighing a solid solute on a balance (giving mass in grams) and measuring solution volume with a volumetric flask (giving volume in liters or milliliters).

How to Calculate Molarity Step by Step

1. Identify the solute and find its molar mass. Use the periodic table to add up the atomic masses of all atoms in the chemical formula. For NaCl: Na (22.99) + Cl (35.45) = 58.44 g/mol.
2. Determine the mass of solute dissolved. This comes from the problem statement or from weighing the solute on a balance.
3. Calculate the moles of solute. Divide the mass by the molar mass: n = mass / molar mass.
4. Determine the total solution volume in liters. Convert milliliters to liters if necessary by dividing by 1,000.
5. Divide moles by volume. M = n / V gives you the molarity.

Dilution Calculations

Dilution is the process of reducing a solution's concentration by adding more solvent. The total moles of solute remain the same before and after dilution — only the volume changes. This principle gives us the dilution equation:

Practical Examples

Other Concentration Units

While molarity is the most common, several other concentration units appear in chemistry. Understanding when to use each one prevents confusion:

Molarity is the standard for most laboratory stoichiometry because it directly connects volume measurements to the number of moles. Molality is preferred when temperature effects must be excluded since it is based on mass rather than volume. Parts per million is the standard for environmental chemistry and water quality testing where concentrations are extremely low.

Preparing Solutions in the Lab

Accurate solution preparation requires careful technique. Follow this protocol for preparing a solution from a solid solute:

1. Calculate the required mass of solute using mass = M x V x molar mass.
2. Weigh the solute accurately on an analytical balance. For solutions requiring high precision, use a balance that reads to 0.0001 g.
3. Transfer the solute to a volumetric flask of the desired final volume. Use a funnel and rinse all residue from the weighing paper and funnel into the flask with small amounts of solvent.
4. Add solvent to approximately 80% of the final volume and swirl or gently mix until the solute is completely dissolved.
5. Bring the volume to the calibration mark by adding solvent dropwise with a wash bottle. Read the meniscus at eye level for accuracy.
6. Mix thoroughly by inverting the stoppered flask several times to ensure homogeneity.

For dilutions from a liquid stock solution, use a pipette to transfer the calculated volume of stock solution to the volumetric flask, then fill to the mark with solvent. Always use calibrated volumetric glassware rather than beakers or graduated cylinders for accurate molarity preparation.

Tips for Working with Molarity

• Always use the molar mass of the actual solute form. If you are using a hydrated salt like CuSO4 · 5H2O, the molar mass includes the water of hydration (249.69 g/mol, not 159.61 g/mol for anhydrous CuSO4). Using the wrong molar mass will give you the wrong molarity.
• Convert milliliters to liters before calculating. The most common arithmetic error in molarity problems is forgetting to convert mL to L. Dividing by 1,000 first prevents this mistake.
• Label all solutions immediately. Every solution container should be labeled with the solute, molarity, date of preparation, and your initials. Unlabeled solutions are a safety hazard and a waste of reagents.
• Account for solute purity. Laboratory-grade reagents are not always 100% pure. If a reagent bottle says 98% purity, you need to weigh 2% more to compensate. Divide the calculated mass by the purity fraction: adjusted mass = calculated mass / 0.98.
• Use volumetric flasks for accuracy. A 250 mL volumetric flask is accurate to within 0.12 mL, while a 250 mL graduated cylinder may only be accurate to within 2 mL. For precise molarity, always use volumetric glassware.

Common Mistakes to Avoid

• Confusing volume of solvent with volume of solution. Molarity uses the total solution volume, not the volume of solvent added. If you add 58.44 g of NaCl to exactly 1,000 mL of water, the total solution volume will be slightly more than 1,000 mL, making the true molarity slightly less than 1.000 M.
• Using mass instead of moles in the formula. Plugging grams directly into M = n/V gives a meaningless number. You must first convert the mass to moles by dividing by the molar mass.
• Forgetting to convert units. If volume is given in milliliters and you need liters for the molarity formula, divide by 1,000. If mass is in milligrams and you need grams, divide by 1,000. Unit conversion errors are the leading cause of incorrect molarity calculations.
• Ignoring hydration in hydrated salts. Many laboratory chemicals are sold as hydrates. Using the anhydrous molar mass instead of the hydrated molar mass for compounds like MgSO4 · 7H2O produces a significant error because the water molecules represent nearly half the mass of each formula unit.
• Rounding too early in multi-step problems. Carry extra significant figures through intermediate steps and only round the final answer to the appropriate number of significant figures. Premature rounding compounds errors across sequential calculations.

Last updated: February 23, 2026