Depending on what is convenient or useful different units of chemical concentration may be used. One common measure is percent by mass. This particular measure is useful when mixing solutions because it is very easy to make a solution with a given concentration. For example, a 20% solution of a solid material may have 20 g of the solid and 80 g of the solvent (usually water). Any other combination with 80/20 solvent/solute ratio will do just as well. Percent by mass allows easy calculation of the mass of solute per gram of solution since it is a simple percentage. For example, a 50 g sample of a 20% solution will have 10 g of the solute. Percent by mass is also useful for measuring very small quantities of dissolved solids. This is employed in giving concentrations for toxic substances such as arsenic or lead in drinking water. Frequently, these measures are converted from percent to parts per million or parts per billion. A percent is a ‘part per hundred’ so a part per million is a percent times 10,000 (10,000 × 100 = 1,000,000).
The most common unit of concentration used in general chemistry labs is molarity, which is the number of moles of solute per liter of solution (mol/L or M). To make a solution with a given molarity the solute needs to be dissolved in a quantity of solvent less than the final volume and mixed thoroughly while it is diluted to the final volume. Since volumes may change as a result of mixing, it is important to add the a final very small amount of solvent to the solution after mixing completely to ensure any volume changes have already taken place. The number of moles of solute to be diluted can be calculated using mass and molar mass or using the concentration of a stock solution with a known molar concentration. In this case the dilution equation (M1V1 = M2V2) makes the process simple. The concentration of the stock solution (M1) will need to have a specific volume (V1) in order for the desired concentration (M2) to be made for the desired final volume (V2). To convert to this concentration unit from mass percent the density of a solution needs to be known. Since percent by mass is based on the total mass of the solution the volume of the solution will need to be calculated for a given mass in order to determine moles per liter of solution. The moles of solute can be calculated based on the mass of solute in a given mass based on the total mass of the solution and the concentration in percent mass.
For applications where the volume of a solution may change, as when temperature causes a liquid to expand or contract in volume, molarity is a poor measure of concentration. Since it is moles per liter of the whole solution it changes if the volume of the solution changes due to thermal expansion upon heating. In these applications, such as the colligative properties of solutions, the concentration unit known as molality is used. Molality is moles of solute per kilogram of solvent. Since it is based on mass it is the same at any temperature. It is simple to convert a mass percent concentration to molality since mass percent means that the mass of solvent for a given mass of solute is known. Simply break down 100 g of the solution into the mass of each component, convert the mass of solute to moles, and divide this by the mass of solvent in kilograms. To convert to molarity, the density must be known.
The final unit of concentration typically used in chemistry courses is the mole fraction. Frequently this is used for mixtures of gases since the volume of a gas is not a fixed quantity and it is difficult to measure the mass of gases. The mole fraction of a component is simply the number of moles of that component divided by the sum of the number of moles of all components in a mixture. For example, a 20% NaCl solution has a mole fraction for NaCl calculated by converting 20 g of NaCl to moles (0.342 mol) and 80 g of water to moles (4.44 mol) and dividing as follows: (0.342)/(0.342 + 4.44) = 0.0715 (or 7.15%). Since mole fraction is often given as a percent it is important to distinguish it from mass percent.
In this brief lab activity you will be given a solution of sodium chloride (NaCl) with a concentration of 20% by mass. Your task will be to calculate its concentration in moles per liter. Using this information you will then mix up 100 mL of a 0.1 M solution of NaCl. You will test your work by making a rough estimate of your chloride concentration using silver nitrate (AgNO3) to precipitate silver chloride (AgCl).
Pre-lab Questions
Answer these questions before carrying out the lab procedure.
The density of a 20% NaCl solution is 1.15 g/mL. What is the concentration of NaCl in mol/L?
Using the given 20% solution as a stock solution, what volume of this solution would be required to make 100 mL of a solution of NaCl with a concentration of 0.10 M?
What are the steps in the procedure for making a solution of known concentration by dilution?
What is the reaction between silver nitrate (AgNO3) and sodium chloride? Give a molecular equation and a net ionic equation.
What volume of 0.20 M AgNO3 solution would be required to react with 1 mL of a 0.1 M NaCl solution?
Describe the appearance of this reaction if it is carried out by adding drops of silver nitrate solution to a sample of sodium chloride solution until the silver nitrate becomes the excess reactant.
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Materials
25 mL 20% NaCl
solution
0.20 M Silver nitrate (AgNO3) solution
100-mL volumetric flask
5 mL Mohr (or serological) pipet
Pipet bulb
Small plastic funnel
Micro-tip Beral pipets
Small test tube
Test tube rack
Centrifuge
Safety
If you choose not to wear safety glasses you are
choosing to sit out the lab.
Silver nitrate is toxic by ingestion. Avoid contact with skin or eyes. It will stain skin a brown color.
Silver chloride contains toxic silver ions and may not be put down the drain or in the trash. Collect this substance as required by your teacher for proper disposal.
Always close containers of chemical supplies when
they are not in use.
Wash hands thoroughly with soap and water before
leaving the laboratory.
Procedure
Be aware of the chemical hazards as you follow this
procedure. If done carefully there is little chance of
harm.
Part I: Making 0.1 M NaCl
Double-check your pre-lab calculations with your lab partner. You should know (1) Conc. in mol/L of the 20% NaCl solution, and (2) The vol. of this solution needed to make 100 mL of a 0.1 M solution.
Obtain your ~25 mL sample of the 20% NaCl solution in a small beaker.
Obtain a 100-mL volumetric flask, a Mohr pipet (graduated, 5 mL capacity), and a pipet bulb.
Measure the required amount of the 20% NaCl solution using the Mohr pipet, as follows:
Ready a clean dry beaker to accept your measured volume of liquid. In this way if you get too much liquid you will not have to wash out the narrow-necked volumetric flask.
Draw the stock solution into the pipet using suction from the bulb up to the zero mark. Make sure the meniscus just touches this mark. Hold it there.
Slowly drain the liquid into the clean beaker. Stop draining when you get to the required volume, which you calculated in the pre-lab.
Drain the remaining stock solution back into your beaker.
Add about 60 - 70 mL distilled water to your beaker with the measured stock solution. Mix.
Use a funnel to transfer your mixture to the 100-mL volumetric flask.
Use a wash bottle full of distilled water to rinse the beaker with a small amount of water. Empty the rinse into the volumetric flask. Repeat the rinsing process 3 or 4 times. This is called a quantitative transfer and it is intended to ensure that all of the intended dissolved substance is transferred to the flask.
Mix the contents of the volumetric flask thoroughly by swirling.
Use the wash bottle or a dropper to add distilled water until the bottom of the meniscus in the narrow neck of the flask just touches the calibration mark.
Part II: Testing your Solution
Now you need to test your work. Use a microtip pipet to add 20 drops of your solution to about 5 mL of distilled water in a small test tube in a rack. Try to keep the drops equal in size. Swirl to mix.
Based on your pre-lab calculations, calculate the number of drops of 0.2 M silver nitrate solution you will need to react with all of the sodium chloride in your test tube.
Obtain a small amount of 0.2 M silver nitrate solution in a clean microtip pipet. Holding the tip a few cm above the surface of the water in your test tube, add the solution drop by drop. The reaction should produce a visible swirl of white precipitate. Keep your eyes close to your test tube to observe this.
Add drops until you have added one less than the number you think will reach stoichiometric equivalence.
The cloudy liquid will no longer allow you to see whether added silver nitrate causes further precipitation. Because of this you will use a centrifuge according to the instruction of your teacher.
In order to operate properly the centrifuge requires that every test tube you place in it has another one with equal mass to balance it out. You can use a test tube filled to the same height with tap water to balance your experimental test tube.
Spin for 5 - 6 minutes to settle the precipitate.
Once centrifugation is complete, test your solution by adding one more drop of silver nitrate. If it causes precipitation, add another drop. If it becomes too cloudy, put it in the centrifuge again for five minutes. Keep track of the total number of drops you add until there is no more precipitation.
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Post-lab
Answer the following questions before you leave the lab to get full credit for your work.
How many drops of 0.2 M AgNO3 did you use? Since there is a 1:1 reltionship between Ag+ and Cl–, and you used 20 drops of your NaCl solution, what concentration does this test say your solution actually has? Assume that all drops have the same volume.
How would adding more than your calculated volume of 20% NaCl solution affect the actual concentration of your intended 0.1 M dilution?
How would adding too much water to your volumetric flask, that is, going above the calibration mark, affect the actual concentration of your intended 0.1 M dilution?
Consider the case that your test solution required 12 drop of 0.2 M silver nitrate solution before precipitation stopped. What would the concentration of the sodium chloride solution be in that case?
How would it affect the concentration of NaCl you calculated based on the silver nitrate test if the conc. of AgNO3 were actually slightly higher than 0.2 M? How would it affect it if the conc. were slightly lower than 0.2 M?
How would your results have been different if the water you used to make your 0.1 M NaCl solution was not distilled but in fact contained dissolved chloride ion?