True solutions are homogeneous mixtures of two or more substances. The substance present in the largest quantity is the solvent, and all other components are called solutes. Homogeneous mixtures form a single phase by eye. (For example, we cannot discern two liquids within a solution, unlike in emulsions—which, therefore, are not true solutions.) In true solutions, the atoms, molecules or ions of the solute are dispersed in the solvent. Special cases are those solutions in which the dispersion of the solute is not molecular (or atomic or ionic) but we cannot visually discern the solute (e.g. micelles). These are not molecular dispersions; yet they do not form a distinct phase. Such solutions are called colloids. The other definition of colloid solutions is that the particle size of the solute is in the range of 1 to 1000 nanometres. The solutions of most proteins fall within this category; i.e. these are colloid solutions, albeit proteins in these solutions are molecular dispersions.
In laboratory practice, mixtures in a liquid state are usually called solutions. In the biochemical laboratory we generally prepare aqueous solutions, i.e. solutions in which the solvent is water. The main reason for this is that we usually work with proteins and other biomolecules that can stably adopt their native conformation in an aqueous environment. We often prepare solutions in organic solvents in order to dissolve smaller compounds (e.g. substrates of enzymes) or in the case of HPLC procedures. Most of the organic solvents used in biochemical experiments are miscible with water. As a semi-quantitative characterisation of solutions, we can distinguish unsaturated, saturated and supersaturated solutions. These solutions contain less, the same, or more than the maximum possible amount of solute, respectively. (These characteristics, together with solubility, are temperature dependent.)
We specify the quantity of components in a mixture quantitatively as their ratio. The precise expression of this can be achieved via concentration units. In liquid-state solutions, it is usually a certain quantity of the solution to which the quantiti(es) of its component(s) are compared, depending on the concentration unit. For instance, in the case of molar concentration (molarity, M, mol/dm3 or mol/L), the most widely used concentration unit in biochemistry, we refer to the number of moles of a given substance in one litre of solution. (For instance, 0.2 M means 0.2 moles of a substance in one litre of solution.) Various percentage forms (%, reference to 100 units of something) are also widely used in biochemistry, although they are not SI units. In the case of weight-by-volume and volume-by-volume percentages (w/v% and v/v%, respectively) we refer to 100 mL of solution, whereas in the case of weight-by-weight percentage (w/w%) we refer to 100 g of solution. In the case of v/v% the volume of the solute, while in the case of w/v% and w/w% the weight of the solute is the basis of reference. (For example, a 15 w/v% solution contains 15 g of solute in 100 mL of solution.) During conversions from or into w/v%, one must also consider the density of the solution (unless it is one g/mL). The density of dilute solutions (up to several %) generally deviates only negligibly from that of the solvent. Therefore, the w/w% and w/v% values of such solutions are practically identical (i.e. 3 w/v% corresponds to ca. 3 w/w%). It is important to note that, for practical reasons, the “biochemical” concentration units mg/mL and μg/mL are also used in the case of protein and other macromolecular solutions in biochemical experiments.
As homogeneity is a basic criterion, it is important to know the solubility of solutes, which is a temperature-dependent property. Information on the solubility of compounds can be obtained from handbook tables (for inorganic compounds) or from the manufacturer (product catalogues of organic compounds). Besides temperature, solubility is influenced by the acidity (pH) of the solvent due to the acid-base character of the solute. Bases tend to dissolve better in acidic solutions, whereas acids do in basic solutions. Therefore, one often needs to change the pH of solutions accordingly. This can be most simply achieved by the addition of an inorganic acid or base (HCl, NaOH, KOH etc.) or by using a buffer system. In both cases, one must be aware of the fact that the prepared solution is a multi-component one, i.e. it contains a pH-setting substance in addition to the substance(s) to be dissolved.
In order to achieve the desired concentration precisely, the measuring range of the devices used (cylinders, pipettes, balances etc.) must be in the range of the quantity to be measured. (For example, a 70-mL solution should not be prepared in a measuring cylinder of a maximal volume of 250 mL, but in one with a 100-mL maximal volume.) As the most widely used concentration units (molarity, w/v% and v/v%) refer the amount of solute in a given volume of the solution (see above), it is the volume of the complete solution that must be set to the desired (calculated) value—thus, one must consider not only the solvent but the solute(s) as well. For instance, when preparing one litre of a 4-M CaCl2 solution, the required amount of CaCl2 must be added not to one litre of water but to much less (e.g. to 500-600 mL, regardless of the condition whether it can dissolve the applied amount of CaCl2), and subsequently complemented with additional water to fill up the volume to one litre. This is the only way to ascertain that the total volume of the solution (CaCl2 and water together) will be one litre. During this procedure, one must take into account that the solute also contributes to the total (final) volume of the solution. Another reason why one must add the solute(s) to an amount of solvent which is less than the final volume might be that one will need to set the pH subsequently in order to enhance solubility and/or to reach a desired pH value. The volume of acid or base solutions used for pH setting is usually not known in advance—and is not negligible. If the solute was initially dissolved in the final desired volume, one would exceed this volume during pH setting (even if the volume of the solute is negligible) and, thus, the solution would not have the desired (calculated) concentration but a smaller (and mostly unknown) one. Thus, according to the correct procedure, one must take into consideration the fact that the final volume will contain the volume of the buffering component.
During mixing the components of a solution, it is often important to be aware of the appropriate and/or safe order of addition. The dilution of concentrated sulphuric acid is a good example of necessary precaution: it is sulphuric acid that must be added to water (and not vice versa) because otherwise water—which boils due to the high heat of solvation—would sprinkle around the hot and corrosive acidic solution, causing serious injury. The appropriate order of dissolving solid solutes is the addition of the solute into the solvent (preferably into less than the final volume aimed)—and not vice versa, pouring solvent onto the solute. This knowledge is of high practical value particularly in the case of dissolving proteins, which have a tendency to stick to the bottom of the dish, thereby making dissolution substantially slower.
Dissolution of substances can be enhanced in various ways including mixing, increasing the temperature, applying sonication or by performing dissolution in the appropriate order (e.g. in the case of the Coomassie solution used to stain proteins in acrylamide gels). Mixing is a commonly used and generally “harmless” method. However, special care must be taken during sonication and especially during heating: one must consider the (heat) stability of the dissolved compound. Proteins are notoriously sensitive molecules; so are a number of simpler organic compounds, too. If we need to set the pH, we must consider the pH sensitivity of the solute. Undesired events and reactions can be prevented by the avoidance of the addition of large quantities or high concentrations of acids or bases. Instead, it is better to use more dilute acid or base solutions and add these in small portions. Sometimes, special requirements must be met such as sterility and/or the sensitivity of the solute to light or oxygen. If the solute is heat sensitive, the solution cannot be sterilised by heat. In these cases, filtration must be applied.
Prepared solutions must appear clear following the complete solution of the solute(s). If this is not the case, the solution must be filtered. (Obviously, one should check if it is not a part of the solute that fell out of the solution.) To this end, one can use traditional tools (filter papers, glass filters) or disposable filters (or filter units) of various materials and pore sizes according to the expected purity of the solution. For example, the solutions used in FPLC, HPLC or in optical devices (photometers, fluorometers) may not contain any floating particles because these are harmful to the device (HPLC, FPLC) or may severely interfere with the measurement.
Solutions are generally stored in glass or plastic containers (bottles, flasks or tubes) that must be tightly closed to prevent changes in concentration due to the slow evaporation of the solvent. (This is much more difficult to achieve in the case of more volatile organic solvents. Here, the use of glass screw-cap vials with Teflon cover is advisable.) For this reason, bakers, Erlenmeyer or other flasks and measuring cylinders (used during preparative procedures) are not useful for this purpose as they do not meet the above criterion of safe storage even in the case of aqueous solutions. Solutions of most of inorganic and organic compounds are generally stored at room temperature. However, in the biochemical laboratory, some solutions must be stored in a refrigerator or a freezer. The latter way of storage is needed to achieve chemical and/or microbiological stability. Unless they are sterilised, solutions of proteins and even those of simple inorganic compounds (e.g. phosphate salts) provide a favourable environment for the growth of bacteria and fungi. (The appearance of these microorganisms is indicated by the increased turbidity and unpleasant smell of the solution. Algae may even appear in distilled water!) Some enzymes (e.g. proteases) may also slowly lose their activity. Cooling—generally below freezing temperature—cannot completely prevent but at least substantially slow down such processes.
Dishes and containers of solutions must be provided with appropriate labels that must remain readable throughout the whole time period of the potential usage of the solution. In the case of solutions containing one or a few different solutes, the label should include the name (or the chemical formula) and the concentration of the solute(s). In many cases, specific names are given to more complex multi-component solutions. These names must be unambiguous at least within the laboratory (e.g. “activation solution”, “sonication buffer”, 10x reaction mixture). Names of many solutions are international if they are widely used during a common laboratory procedure. Such typical names include those of reagents (e.g. the Bradford reagent used to measure protein concentration) or culture media (e.g. LB or 2YT used to grow bacteria). The time of preparation of a solution may be of importance and thus it should be included in the label—the name of the person who prepared the solution is also an option.