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Chapter 4: Problem 111

A solution is prepared by dissolving \(0.6706 \mathrm{~g}\) oxalic acid \(\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) in enough water to make \(100.0 \mathrm{~mL}\) of solution. \(\mathrm{A}\) \(10.00-\mathrm{mL}\) aliquot (portion) of this solution is then diluted to a final volume of \(250.0 \mathrm{~mL}\). What is the final molarity of the oxalic acid solution?

Short Answer

Expert verified
The final molarity of the oxalic acid solution after dilution is 0.00298 M.

Step by step solution

01

Calculate the initial moles of oxalic acid

To find the initial moles of oxalic acid, use the given mass and the molar mass of oxalic acid. The molecular formula of oxalic acid is \(H_2C_2O_4\), so its molar mass is: Molar mass of oxalic acid = 2 (H) + 2 (C) + 4 (O) = 2(1.01) + 2(12.01) + 4(16.00) = 90.04 g/mol Now, use the mass of oxalic acid to calculate the moles: Moles of oxalic acid = mass of oxalic acid / molar mass of oxalic acid = 0.6706 g / 90.04 g/mol = \(7.45 \times 10^{-3}\) mol
02

Calculate the initial molarity of the solution

Now that we have the moles of oxalic acid, we can find the initial molarity using the volume of the initial solution. Molarity (M) = moles of oxalic acid / volume in L = \(7.45 \times 10^{-3}\text{ mol} / 0.100\text{ L}\) = 0.0745 M
03

Use the dilution concept to find the final molarity

When a portion of the initial solution is diluted, the moles of solute in both solutions remain the same. So, we can use the equation: M1V1 = M2V2 where M1 is the initial molarity, V1 is the aliquot volume, M2 is the final molarity, and V2 is the final volume. We know M1, V1, and V2, so we can find M2: \(0.0745 \text{ M} \times 0.010 \text{ L} = M2 \times 0.250 \text{ L}\) Now, solve for M2: M2 = \((0.0745 \times 0.010) / 0.250 \text{ M}\) = 0.00298 M
04

Final molarity of the oxalic acid solution

The final molarity of the oxalic acid solution after dilution is 0.00298 M.

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Most popular questions from this chapter

Give an example how each of the following insoluble ionic compounds could be produced using a precipitation reaction. Write the balanced formula equation for each reaction. a. \(\mathrm{Fe}(\mathrm{OH})_{3}(s)\) b. \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}(s)\) c. \(\mathrm{PbSO}_{4}(s)\) d. \(\mathrm{BaCrO}_{4}(s)\)

A \(1.42-\mathrm{g}\) sample of a pure compound, with formula \(\mathrm{M}_{2} \mathrm{SO}_{4}\), was dissolved in water and treated with an excess of aqueous calcium chloride, resulting in the precipitation of all the sulfate ions as calcium sulfate. The precipitate was collected, dried, and found to weigh \(1.36 \mathrm{~g}\). Determine the atomic mass of \(\mathrm{M}\), and identify \(\mathrm{M}\).

What volume of \(0.0200 M\) calcium hydroxide is required to neutralize \(35.00 \mathrm{~mL}\) of \(0.0500 \mathrm{M}\) nitric acid?

A mixture contains only \(\mathrm{NaCl}\) and \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3} .\) A \(1.45-\mathrm{g}\) sample of the mixture is dissolved in water and an excess of \(\mathrm{NaOH}\) is added, producing a precipitate of \(\mathrm{Al}(\mathrm{OH})_{3}\). The precipitate is filtered, dried, and weighed. The mass of the precipitate is \(0.107 \mathrm{~g}\). What is the mass percent of \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) in the sample?

The blood alcohol \(\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{OH}\right)\) level can be determined by titrating a sample of blood plasma with an acidic potassium dichromate solution, resulting in the production of \(\mathrm{Cr}^{3+}(a q)\) and carbon dioxide. The reaction can be monitored because the dichromate ion \(\left(\mathrm{Cr}_{2} \mathrm{O}_{7}{ }^{2-}\right)\) is orange in solution, and the \(\mathrm{Cr}^{3+}\) ion is green. The balanced equation is \(16 \mathrm{H}^{+}(a q)+2 \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{OH}(a q) \longrightarrow\) \(4 \mathrm{Cr}^{3+}(a q)+2 \mathrm{CO}_{2}(g)+11 \mathrm{H}_{2} \mathrm{O}(t)\) This reaction is an oxidation-reduction reaction. What species is reduced, and what species is oxidized? How many electrons are transferred in the balanced equation above?
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