Page 211 - Demo

Percent Solutions
Mass percent solutions are defined based on the grams of solute per 100 grams of solution.
Example: 20 g of sodium chloride in 100 g of solution is a 20% by mass solution.
Volume percent solutions are defined as milliliters of solute per 100 mL of solution.
Example: 10 mL of ethyl alcohol plus 90 mL of H2O (making approx. 100 mL of solution) is a 10% by volume solution.
Mass-volume percent solutions are also very common. These solutions are indicated by w/v% and are defined as the grams of solute per 100 milliliters of solution.
Example: 1 g of phenolphthalein in 100 mL of 95% ethyl alcohol is a 1 w/v% solution.
Conversion Between Percent Solutions
You may wish to convert mass percent to volume percent or vice versa. If so, follow this procedure:
A 10% by mass solution of ethyl alcohol in water contains 10 g of ethyl alcohol and 90 g of water.
1. The formula for determining the volume of the component (ethyl alcohol in our example) is:
3. Determine the percent by volume by dividing the volume of the component by the volume of the solution.
mass of ethyl alcohol Volume = ——————————
density of ethyl alcohol
2. Determine the volume of the total solution by dividing the mass of the solution by the density of the solution.
Let’s solve
1. Mass of Density
Volume
Volume
2. Mass of Density
Volume
3. Volume Percent Reverse
percent.
Calculating Molarity from Percent Solutions
To determine the molarity of a mass percent solution, the density of the solution is required. Use the following procedure:
1. Determine the mass of solution by multiplying the volume of the solution by the density of the solution.
mass = volume x density
2. Determine concentration in percent by mass of the solute in
solution. Change to the decimal equivalent.
3. Calculate the molar mass of the compound, MM.
4. Multiply mass (step 1) by mass % (step 2) and divide by molecular mass (step 3) to find the number of moles present in the whole solution.
5. Divide the number of moles (step 4) by the volume in liters of the solution to find the molarity of the solution.
Example: Determine molarity of 37.2% hydrochloric acid (density 1.19 g/mL).
1. Mass of solution = 1,000 mL x 1.19 g/mL = 1,190 g
2. Mass % = 37.2 % = 0.372
3. Molar mass of hydrochloric acid = 36.4 g/mol
1, 2, and 3 above as follows:
ethyl alcohol = 10 g (given)
of ethyl alcohol = 0.794 g/mL (from handbook)
mass = ——— density
10 g
of ethyl alcohol = ————— = 12.6 mL
0.794 g/mL solution = 100 g (given)
of solution (10% ethyl alcohol) = 0.983 g/mL (from
handbook)
100 g
of solution = ————— = 101.8 mL* 0.983 g/mL
percent of solution
volume of ethyl alcohol 12.6
= —————————— = ——— = 12.4%
total volume of solution 101.8
the procedure to convert volume percent to mass
4. mass x mass % 1,190 g x 0.372 ———————— = ———————— = 12.1 moles
MMHCl 36.4 g/mol
5. Molarity = moles/liters = 12.1 moles/1 liter = 12.1 M
* The volume percent statement generally is accurate but the volume percent is not always calculated directly from the volumes of the mixed ingredients because the final volume may not equal the sum of the separate volumes. In our solution (No. 2 above) note that if the alcohol volume (12.6 mL) is added to the water volume (90 mL), the final volume is less than 102.6 mL.
Normality: A concentration unit (N); defined as the number of equivalents of solute per liter of solution. (e.g., 1 M H2SO4 = 2 N H2SO4)
Saturated Solution: A solution that contains the maximum amount of a particular solute that will dissolve at that temperature.
Solute: The substance which is dissolved, or has gone into solution (typically a solid).
Solution: A uniform homogeneous mixture of two or more substances. The individual substances may be present in varying amounts.
Solvent: The substance which does the dissolving (typically a liquid, such as water or alcohol). Must be greater than 50% of the solution.
Standard Solution: A very precise solution, usually to 3–4 significant figures, used in quantitative analysis or an analytical procedure.
Supersaturated Solution: A solution that contains more solute than equilibrium conditions allow; it is unstable and the solute may precipitate upon slight agitation or addition of a single crystal.
“Your Safer Source for Science” Laboratory Solution Preparation Basic Concepts of Preparing Solutions, continued
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