How to Accurately Prepare Laboratory Solutions

Preparing a laboratory solution means dissolving a measured amount of solute in a solvent to obtain a known concentration, either by weighing a solid and diluting it to an exact volume or by diluting a more concentrated stock solution. In brief: to make 250.0 mL of 0.100 mol/L sodium chloride, weigh 1.461 g of NaCl, dissolve it in less than 250 mL of water, transfer it to a 250.0 mL volumetric flask and add water to the mark.

The accuracy of almost every experiment in solution chemistry depends on how well the solutions were prepared. A small mistake in the preparation shows up as an error in all later results. This article explains the common ways of expressing concentration, how to prepare a solution from a solid and from a stock solution, explains percent solutions and serial dilutions, includes worked calculations and describes the habits that give accurate results.

How is concentration expressed?

How do you prepare a solution from a solid?

1. Calculate the mass of solute needed: mass = molarity × volume (in L) × molar mass.
2. Weigh the solute on a balance, using a weighing boat. Record the actual mass, which may be slightly different from the target.
3. Dissolve the solute in a beaker with about half to two thirds of the final volume of solvent. Stir until it is completely dissolved. Warm gently if necessary, but let the solution cool to room temperature afterward.
4. Transfer the solution into a volumetric flask of the required volume using a funnel. Rinse the beaker, the stirring rod and the funnel with small portions of solvent, and add the rinsings to the flask.
5. Dilute to the mark. Add solvent until the level is near the mark, then add drop by drop until the bottom of the meniscus touches the mark at eye level.
6. Mix. Stopper the flask and invert it at least ten times.
7. Label the container with the name, concentration, date and your initials.

How do you prepare a solution by dilution?

Diluting a stock solution does not change the number of moles of solute, so C₁V₁ = C₂V₂, where 1 is the stock and 2 is the diluted solution. Solve for the volume of stock needed, V₁ = C₂V₂ ÷ C₁. Measure that volume with a pipette or burette, transfer it into a volumetric flask partly filled with solvent, and dilute to the mark. When diluting a concentrated acid, always add the acid to the water.

Worked examples

1. From a solid. Prepare 250.0 mL of 0.100 mol/L NaCl. Moles = 0.100 × 0.2500 = 0.0250 mol. Mass = 0.0250 × 58.44 = 1.461 g.
2. Adjusting for the actual mass. The student weighs 1.468 g instead of 1.461 g. The real concentration is (1.468 ÷ 58.44) ÷ 0.2500 = 0.02512 ÷ 0.2500 = 0.1005 mol/L. Use the real value in later calculations, and do not try to remove salt.
3. From a hydrate. Prepare 500.0 mL of 0.0500 mol/L copper(II) sulfate from CuSO₄·5H₂O (249.7 g/mol). Moles = 0.0500 × 0.5000 = 0.0250 mol. Mass = 0.0250 × 249.7 = 6.24 g. The molar mass of the hydrate must be used, since the water is part of the weighed mass.
4. By dilution. Prepare 100.0 mL of 0.250 mol/L HCl from 12.0 mol/L stock. V₁ = (0.250 × 100.0) ÷ 12.0 = 2.08 mL of stock, diluted to 100.0 mL.
5. Percent solution. A 0.9% (m/v) saline solution has 0.9 g of NaCl in each 100 mL. For 500 mL: 0.9 × 500 ÷ 100 = 4.5 g of NaCl, dissolved and diluted to 500 mL.
6. Mass percent. To make 200 g of a 5.0% (m/m) sugar solution: sugar = 0.050 × 200 = 10 g, and water = 190 g.
7. Parts per million. A solution with 2.5 mg of fluoride ions in 500 mL has 2.5 ÷ 0.500 = 5.0 mg/L = 5.0 ppm.

What is a serial dilution?

A serial dilution makes a series of solutions, each diluted by the same factor from the previous one. It is used to produce a set of standards of known concentrations, for example for calibration curves, and to dilute very concentrated samples in small steps. If a 1.00 mol/L solution is diluted tenfold three times (1.00 mL into 9.00 mL each time), the concentrations are 0.100, 0.0100 and 0.00100 mol/L.

How do you store and label solutions?

Store solutions in clean, labeled bottles of a suitable material. Some solutions are unstable: sodium hydroxide solutions absorb carbon dioxide from the air and should be kept tightly closed in plastic bottles, silver nitrate solutions are kept in dark bottles to avoid decomposition by light, and dilute solutions of some compounds should be made fresh. Always record the date and the preparer, and discard solutions after their shelf life or when they look cloudy or discolored.

How did standard solutions develop?

The use of solutions of known concentration for analysis, called volumetric analysis, grew in the 19th century with the work of Gay-Lussac and Mohr. The concept of molar concentration became widespread with the development of the mole concept in the early 20th century. In 1909 the Danish chemist Søren Sørensen introduced the pH scale while working on the effect of acidity on enzymes, which made accurately prepared solutions and buffers even more important in biochemistry.

Which mistakes are common?

Mistake 1: Dissolving the solute directly in the volumetric flask to the mark

The solute adds volume. Dissolve first, then fill to the mark.

Mistake 2: Using the molar mass of the anhydrous salt for a hydrate

Use the molar mass of the substance as it is in the bottle, including the water of crystallization.

Mistake 3: Diluting by adding the stock volume to the final volume of water

Adding 10.0 mL of stock to 100.0 mL of water gives about 110 mL, not 100.0 mL. Dilute to the mark.

Mistake 4: Not mixing after adding solvent

Concentration in the lower part of the flask can be different. Invert the flask several times.