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# MOLARITY / CONCENTRATION - North Carolina School of … - mol l to g ml

MOLARITY / CONCENTRATION

NAME:__________________________________________________ PERIOD:___________

Prelab

1. Calculate the mass of compound required to prepare 100.00 ml of a 1.40x10-2M KMnO4. Show your work.

2. If 2.50g of NaOH are dissolved in 100.00mL of water, what is the molarity of the solution? Show your work.

3. If 10.00mL of the solution in question 2 are diluted to 100.00mL, what will be the molarity of the dilute solution? Show your work.

4. If 10.00mL of the solution in question 2 are combined with to 25.00mL of the solution from question 3 in a 100mL volumetric flask and diluted to 100.00mL, what will be the molarity of the final solution? Show your work.

MOLARITY / CONCENTRATION

Objective: There are three objectives to this lab:
1) To learn how to prepare a solution which is a specific molarity (M).
2) To carry out a dilution and calculate its molarity.
3) To combine specific volumes of the two above solutions and calculate the mixture’s molarity.

Introduction: Many chemical reactions done in the laboratory involve some type of aqueous solution. The concentration of the solute in these solutions is important. Since the stoichiometry of a reaction involves moles and mole ratios, it would helpful if the concentration unit also involved moles. A convenient way to express the concentration of a solution is molarity (M). Molarity is defined as moles of solute dissolved per liter of solution produced.

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The convenience of molarity is that the volume of a given solution is easily measured. For a solution with a known molarity, the moles of the solute in the solution are easily calculated. For example, 250 ml of a 1.00M KCl solution has 0.250 moles of the solute.

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Once the moles are calculated, the mass of KCl required to prepare the solution can be calculated.

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If a known amount of solid is dissolved and diluted to a final volume, the molarity of the solution can be calculated. If you dissolve 0.7500 g of NaCl in water and dilute it to a final volume of 100.00mL, you would calculate the molarity as follows:

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Because molarity is defined in terms of the volume of solution, special volumetric glassware, volumetric flasks, must be used for very accurate work. This glassware is carefully calibrated to contain a specific volume: 1000.00 ml, 500.00 ml, 10.00 ml, etc. The volume that a liquid occupies changes with temperature. As a solution’s temperature is increased the liquid expands and as a result occupies a greater volume. As the solution’s temperature is decreased the liquid contracts and as a result occupies a lesser volume. Therefore, volumetric glassware is calibrated at a particular temperature (often 20oC), which is approximately room temperature. Volumetric glassware should not, therefore be used to measure hot or cold liquids.
There are additional advantages to using molarity for concentration. Since the solution is diluted to the final volume, the mass and volume of the solvent used does not have to be known. This means that the concentration is independent of the solvent used. A 1.00M solution in water has the same concentration as a 1.00M solution in ethanol. It would also be difficult to mass 1x10-6 mole of a pure compound even on an analytical balance. One can relatively easily prepare a
1x10-3M solution and transfer 1mL of the solution to a reaction. This would produce 1x10-6 mole,
(1x10-3M)(1x10-3L) = 1x10-6 mole. Very small numbers of moles and grams of compound can be used in a reaction by using dilute solutions.
Once a solution is prepared, it is often economical and convenient to prepare a new solution by diluting a known volume of a solution of a known concentration to a new final volume with the solvent. The moles in the two solutions will be equal.

Moles in the concentrated solution = Moles in the dilute solution

Moles = (M)(L)

(Molarityconc. soln.)(Vconc. soln.) = (Molaritydilute soln.)(Vdilute soln.)

If 10.00mL of 6.00M HCl are diluted to 250.0mL, the final concentration will be 0.240M.

(6.00M)(1.000x10-2L) = (xM)(0.2500L)

xM = 0.240M

New solutions can also be prepared by mixing several solutions. If the solutions are thoroughly mixed, the total numbers of moles are dispersed throughout the total volume. The final molarity is given by the expression

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If 50.00mL of 0.100M HCl are added to 100.00mL of 0.400M HCl and the mixture is diluted to a final volume of 500.00mL, the final molarity will be

Solution 1 (0.100M HCl)(0.05000L) = 5.00x10-3 mol HCl

Solution 2 (0.400M HCl)(0.1000L) = 4.00x10-2 mol HCl

Total moles = 4.50x10-2 mol HCl

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Procedure: You will prepare three solutions during the course of this laboratory: 1) a stock solution, 2) a dilution of the stock solution, and 3) a mixture of the stock and dilution. The solutions to be prepared are as follows:

Stock (solution 1)
1. Mass between 0.0750g and 0.1000g of KMnO4 in a small plastic weighing boat on an analytical balance. This is a very small amount of compound. Add it carefully to the weighing boat. Clean up any spilled compound immediately. This compound is an oxidant. Avoid skin contact.

Dilution (solution 2) Prepare 100.00 mL of a dilute solution of KMnO4 from the stock solution as follows:
1. Transfer 5.00mL of the stock solution into a 100.00mL volumetric flask using a 5.00mL volumetric pipet.
2. Fill the volumetric flask to the 100.00 mark with deionized water.
3. Stopper the flask and mix the solution again. Remember to keep your fingers on the stopper. Label the flask “dilution” and set it aside for later.

Mixture (solution 3) Prepare 100.00 mL of a mixture of the stock solution and the dilution as follows.
1. Transfer 5.00mL of the stock solution into a 100.00mL volumetric flask using a 5.00 mL volumetric pipet.
2. Transfer 25.00mL of the dilute solution into the same 100.00mL volumetric flask using a 25.00mL volumetric pipet.
3. Fill the volumetric flask to the 100.00 mark with deionized water.
4. Stopper the flask and mix the solution. Remember to keep your fingers on the stopper. Label the flask “mixture”.

Record any visual observations that relate to concentration differences between the solutions. When you have completed the observations, take your pipets, volumetric flasks, and stoppers to the sink and rinse them with hot water. When they are clean, dry the outside with paper towels and return them to your workstation.

Calculations:
1. Calculate the molarity of each solution based on the mass of the compound or the volume of the solution used.

MOLARITY / CONCENTRATION

NAME____________________________________________________PERIOD__________

LAB PARTNER____________________________________________COURSE_________

Data Table
Stock Solution:
|Mass of compound | |
|Molar mass of compound | |
|Moles of compound | |
|Volume of stock solution in liters | |
|Molarity of the stock solution | |

Dilution:
|Volume of stock solution pipetted in mL | |
|Volume of stock solution pipetted in L | |
|Moles of compound pipetted | |
|Volume of dilute solution in liters | |
|Molarity of the dilute solution | |

Mixture
|Volume of stock solution pipetted in mL | |
|Volume of stock solution pipetted in L | |
|Moles of compound pipetted | |
|Volume of dilute solution pipetted in mL | |
|Volume of dilute solution pipetted in L | |
|Moles of compound pipetted | |
|Total moles of compound pipetted | |
|Volume of mixed solution in liters | |
|Molarity of the mixed solution | |

Calculations: Show your calculation of the molarity of each solution.