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Chapter 4: Problem 60

Describe the steps involved in preparing a solution of known molar concentration using a volumetric flask.

Short Answer

Expert verified
To prepare a molar solution using a volumetric flask, begin by calculating the required amount of solute in grams. Measure this amount, add it to a clean and dry volumetric flask, then add solvent up to the calibration line. Mix thoroughly to ensure a homogeneous solution.

Step by step solution

01

Calculation of the Required Amount of Solute

Calculate the required amount of solute by multiplying the desired molar concentration by the volume of the solution. The result will give the amount of solute required in moles. Depending on the particular exercise, this amount may need to be converted from moles to grams by multiplying by the molar mass. This is done using the formula: mass (g) = volume (L) * concentration (M) * molecular weight (g/mol).
02

Measurement of Solute

Using a suitable weighing balance, measure the calculated amount of solute. If the solute is a liquid, it should be measured using a pipette or another suitable liquid measure.
03

Addition of Solute to Volumetric Flask

Carefully transfer the solute into a clean and dry volumetric flask. If the solute is a solid, use a carefully folded piece of clean paper to guide it into the flask.
04

Addition of Solvent to Volumetric Flask

Slowly add the solvent to the flask, rinsing any remaining solute into the flask in the process. Continue adding solvent until the bottom of the meniscus lies on the etched calibration line on the neck of the flask.
05

Mixing of Solute and Solvent

Finally, stopper the volumetric flask and invert it several times to ensure the solute is completely dissolved and evenly distributed, forming a homogeneous solution.

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Most popular questions from this chapter

Calculate the volume in milliliters of a \(1.420 \mathrm{M}\) \(\mathrm{NaOH}\) solution required to titrate the following solutions. (a) \(25.00 \mathrm{~mL}\) of a \(2.430 \mathrm{M} \mathrm{HCl}\) solution (b) \(25.00 \mathrm{~mL}\) of a \(4.500 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution (c) \(25.00 \mathrm{~mL}\) of a \(1.500 \mathrm{M} \mathrm{H}_{3} \mathrm{PO}_{4}\) solution

Oxalic acid \(\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) is present in many plants and vegetables. If \(24.0 \mathrm{~mL}\) of \(0.0100 \mathrm{M} \mathrm{KMnO}_{4}\) solution is needed to titrate \(1.00 \mathrm{~g}\) of a sample of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) to the equivalence point, what is the percent by mass of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) in the sample? The net ionic equation is \(2 \mathrm{MnO}_{4}^{-}+16 \mathrm{H}^{+}+5 \mathrm{C}_{2} \mathrm{O}_{4}^{2-} \longrightarrow\) \(2 \mathrm{Mn}^{2+}+10 \mathrm{CO}_{2}+8 \mathrm{H}_{2} \mathrm{O}\)

The recommended procedure for preparing a very dilute solution is not to weigh out a very small mass or measure a very small volume of a stock solution. Instead, it is done by a series of dilutions. A sample of \(0.8214 \mathrm{~g}\) of \(\mathrm{KMnO}_{4}\) was dissolved in water and made up to the volume in a 500-mL volumetric flask. A \(2.000-\mathrm{mL}\) sample of this solution was transferred to a \(1000-\mathrm{mL}\) volumetric flask and diluted to the mark with water. Next, \(10.00 \mathrm{~mL}\) of the diluted solution were transferred to a \(250-\mathrm{mL}\) flask and diluted to the mark with water. (a) Calculate the concentration (in molarity) of the final solution. (b) Calculate the mass of \(\mathrm{KMnO}_{4}\) needed to directly prepare the final solution.

A sample of \(0.6760 \mathrm{~g}\) of an unknown compound containing barium ions \(\left(\mathrm{Ba}^{2+}\right)\) is dissolved in water and treated with an excess of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\). If the mass of the \(\mathrm{BaSO}_{4}\) precipitate formed is \(0.4105 \mathrm{~g},\) what is the percent by mass of Ba in the original unknown compound?

Describe how you would prepare \(250 \mathrm{~mL}\) of a \(0.707 M \mathrm{NaNO}_{3}\) solution.
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