Preparing solutions

A solution is a system in which one or more solutes (which can be solid, liquid, or gas) is dissolved in a liquid solvent. One common lab practice that all chemists need to become very familiar with is the preparation of solutions. There are many different ways to prepare solutions, and this page will cover the most common methods in detail.

Basic concepts

Solutions consist of one or more solutes dissolved in a solvent. For the purposes of this tutorial, the solutes can be solid, liquid, or gas, and the solvent is always a liquid, although technically solutions can be composed of solutes and solvents in any state. Solutions are defined by their concentration, which depending on the purpose of the solution, can be measured in multiple ways.

There is almost always a maximum amount of solute that can be dissolved in a given solvent at a given temperature. Once this amount has been dissolved, the solution is considered to be saturated. The value of the concentration of the saturated solution is known as the solute's solubility in the given solvent at the given temperature. Increasing temperature will usually increase the solubility of solids and decrease the solubility of gases, although there are a few solid solutes that are more soluble at lower temperatures, such as lithium carbonate or cerium(III) sulfate; these compounds may be said to have inverse solubility curves. There are some liquid-liquid systems that are considered to be miscible, meaning that the components are soluble within one another in all proportions. An example of this is ethanol and water.

Units of concentration

There are many different ways to measure concentration. This section will cover the ones most commonly used in the lab.

Molarity

One of the most common concentration units is molarity. Molarity is measured in moles of solvent per liters of solution. It can be abbreviated as mol/L or simply M. Molar solutions are very useful in the lab, as it is very easy to determine the volume of solution needed based on the stoichiometry. For instance, if 0.2 mol of a substance is needed, 100 mL of a 2 M solution could be easily measured out.

Preparation of a Molar Solution

• First, determine the molar mass of the solute. This can be found in a reference book or online. If it is a hydrate, make sure to take the mass of the water of hydration into account. (For example, copper(II) sulfate pentahydrate, 249.686 g/mol)
• Next, consider what a reasonable concentration of the solution would be. Make sure that the solute is fully soluble at the desired concentration. Otherwise, it will not fully dissolve. (0.8 M)
• Decide what volume of solution is needed. This may be constrained by the available equipment. It is best to prepare molar solutions in volumetric flasks, although graduated cylinders will also work if volumetric flasks are not available. (250 mL)
• Weigh out the needed mass of solute. If it is a solid, crystalline compound, it would be prudent to crush it into a powdered form with a mortar and pestle prior to weighing it so that it will dissolve faster. (0.25 L * 0.8 M * 249.686 g/mol = 49.937 g)
• Add the solute to the proper size of volumetric flask or graduated cylinder before adding the solvent using a powder funnel to prevent spills. (250 mL)
• Rinse the funnel and the weighing boat with a small amount of the solvent and add it to the flask. Then fill the flask about halfway with the solvent and swirl.
• Carefully continue adding the solvent until the bottom of the meniscus is touching the graduate on the flask.
• Stopper the flask and repeatedly invert it until the solute is completely dissolved. Make sure to hold on to the stopper while doing so.
• Transfer the solution to a bottle for storage and label it with the name of the solute and the concentration. (CuSO₄, 0.8 M)

Molality

Another concentration unit that is fairly common to see, although not as common as molarity, is known as molality. It is a measure of the moles of solute per kilograms of solvent, not liters of solution. Because of this, molal solutions are less convenient to work with than molar solutions, as the amount of solution needed can not be determined immediately from the volume. There are some advantages though. For one, molality is constant with temperature and pressure, unlike molarity, since the mass of substances does not change with temperature and pressure, while the volume does. Also, masses can generally be measured more accurately than volumes, so molality has the potential to be more precise. It can be abbreviated as mol/kg or simply m.

Preparation of a Molal Solution

• First, determine the molar mass of the solute. This can be found in a reference book or online. If it is a hydrate, make sure to take the mass of the water of hydration into account. (For example, sodium bromide, 102.89 g/mol, aqueous)
• Next, decide on what the molality of the solution will be, and what mass of solvent will be used. Check a reference to make sure that the compound is soluble at the desire concentration. For dilute aqueous solutions, values of molarity and molality are usually very close. (1.5 m, 500 g H₂O)
• Weigh out the needed mass of solute. If it is a solid, crystalline compound, it would be prudent to crush it into a powdered form with a mortar and pestle prior to weighing it so that it will dissolve faster. (1.5 m * 0.5 kg * 102.89 g/mol = 77.168 g)
• Weight out the needed mass of solvent. If the solvent is water, it may be easier to estimate the density of water as 1 g/mL instead of weighing the water, although this isn't as precise. (500 g or ~500 mL)
• Add the solute to a beaker before adding the solvent.
• Rinse the solute's weighing boat with a small amount of the solvent and add it to the beaker. Then add the rest of the solvent and stir until it is fully dissolved.
• Transfer the solution to a bottle for storage and label it with the name of the solute and the concentration. (NaBr, 1.5 m)