The following pages provide you with a summary of solution types. Hold on to these pages during your chemistry, biology and other science courses. They will come in handy.
Work in the fields of synthesis, quantitative analysis, and qualitative analysis all depend on the ability to prepare solutions of known concentration. Virtually all aspects of life are affected, from food preparation to medical diagnostics to manufacturing to the extraction of compounds from exotic plants. There are several aspects of solution preparation that are used and each method of preparation followed by a calculation and exercise will be performed. Solutions are made up of two components. The solvent is the compound, usually present in large excess, in which the solute is dissolved. The solvent is usually a liquid at room temper- ature. Water is the most commonly used solvent. The solvent itself may be a solution, say 0.1 M hydrochloric acid(HCl). A solution such as 1.0M barium chloride in 0.1M HCl can also be thought of as a solvent(water) with two solutes (HCl and BaCl2). The units of concentration vary widely. The most common unit used by chemists is molarity (M, moles per liter of solution). Others include molality (m, moles of solute per kg of solvent), normality((N, equivalents per liter), and in the medical field, mg per deciliter (deciliter = 0.1 L). Less commonly in chemistry but commonly used in commercial packaging are percents - for solid solutes, weight:volume (w/v); and for liquids, volume:volume(v/v). For example, a 10% (w/v) sodium bicarbonate solution contains 10g of sodium bicarbonate per 100 mL of solution, whereas a 5% (v/v) alcohol solution contains 5 mL of alcohol per 100 mL of solution. In expressing the concentrations of very dilute solutions, percentage compositions become awkward to use because of the number of zeros needed to place the decimal point. The concentration is more conveniently expressed in parts per million (ppm). This term is defined as the
for solid mixtures: ppm = mass of solute in g x 103 mg x 103 g mixture = mass of solute in g x 1,000,000 = mass of mixture in g g solute 1 kg mass of mixture in g
ppm = mg solute kg mixture
for solutions (assuming a density of 1g/1mL): ppm = mass of solute in g x 103 mg x 103 mL = g ∙ 1,000,000 mg ∙ mL = volume of solution in mL g 1L mL g ∙ L
ppm = mg L
Remember that 1 g of water occupies 1mL. For aqueous solutions, it is easiest to think of ppm as mg/L.
Preparation of Solution by Dissolving a Solid
Most of the solutions you will prepare will have the concentration expressed in units of molarity. When the desired volume and concentration of a solution to be made up from a solid is known, the required mass of compound is calculated as follows:
No. of moles required = M x V = (mol/L)(L) (1)
mass = moles x MM = (mol)(g/mol) (2)
Combining these two equations: mass needed = M x V x MM = (mol/L)(L)(g/mol) (3)
Example 1: You are asked to make up 100 mL of 0.25M KOH. M = 0.25 mol/L, V = 0.10 L, MM of KOH = 56.11 g/mol
mass of KOH needed = (0.25 mol/L)(0.10 L)(56.11 g/mol) = 1.40275 g KOH = 1.4 g KOH (2 Sig. Figs.) This mass is emptied into a volumetric flask and filled with enough water to dissolve. Then water is carefully added to bring the solution up to the 100 mL calibration line, with occasional swirling during, and final capping and inversion afterward.
Dilution
When a solution is to be made up by diluting another solution, the equation is slightly different. We must determine the number of moles of solute wanted in the solution. We then calculate what volume of the more
