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Chapter 10: Problem 12

Describe how you would prepare \(1.00 \mathrm{~L}\) of \(0.750 \mathrm{M}\) barium hydroxide solution starting with (a) solid barium hydroxide. (b) \(6.00 \mathrm{M}\) barium hydroxide solution.

Short Answer

Expert verified
Answer: (a) To prepare the solution using solid barium hydroxide, weigh out 128.50 g of solid Ba(OH)_2 and dissolve it in roughly 800 mL of deionized water. Stir until the solid is completely dissolved. Transfer the solution to a 1.00 L volumetric flask and carefully add deionized water until the 1.00 L mark is reached, mixing thoroughly. (b) To prepare the solution using a 6.00 M Ba(OH)_2 solution, measure 125 mL of the 6.00 M Ba(OH)_2 solution and transfer it to a 1.00 L volumetric flask. Add deionized water to the flask until the 1.00 L mark is reached, mixing thoroughly.

Step by step solution

01

Calculate the moles of barium hydroxide required

To calculate the moles of barium hydroxide needed, use the formula: moles = molarity × volume, where molarity is 0.750 M and volume is 1.00 L (converted to liters). moles = 0.750 M * 1.00 L moles = 0.750 mol
02

Convert moles to grams

Now, find the molar mass of Ba(OH)_2: Ba(OH)_2 = 137.33 g/mol (Ba) + 2 * (15.999 g/mol (O) + 1.008 g/mol (H)) Ba(OH)_2 ≈ 171.34 g/mol Next, convert the moles of Ba(OH)_2 to grams: 0.750 mol × 171.34 g/mol ≈ 128.50 g
03

Prepare the solution

Weigh out 128.50 g of solid Ba(OH)_2 and dissolve it in roughly 800 mL of deionized water. Stir the solution until the solid is completely dissolved. Transfer the solution to a 1.00 L volumetric flask and carefully add deionized water until the 1.00 L mark is reached, mixing thoroughly. Your 0.750 M Ba(OH)_2 solution is now prepared. (b) Preparing a solution using a 6.00 M barium hydroxide solution:
04

Determine the volume of 6.00 M Ba(OH)_2 solution needed

Use the dilution equation: M1 * V1 = M2 * V2 Where M1 and V1 are the molarity and volume, respectively, of the more concentrated solution (6.00 M Ba(OH)_2), and M2 and V2 are the molarity and volume of the target solution (0.750 M Ba(OH)_2 and 1.00 L). Rearrange the equation to solve for V1: V1 = (M2 * V2) / M1 V1 = (0.750 M * 1.00 L) / 6.00 M V1 ≈ 0.125 L = 125 mL
05

Prepare the solution

Measure 125 mL of the 6.00 M Ba(OH)_2 solution using a graduated cylinder or pipette. Transfer the measured solution to a 1.00 L volumetric flask. Add deionized water to the flask until the 1.00 L mark is reached, mixing thoroughly. Your 0.750 M Ba(OH)_2 solution is now prepared.

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Most popular questions from this chapter

Explain why (a) the freezing point of \(0.10 \mathrm{~m} \mathrm{CaCl}_{2}\) is lower than the freezing point of \(0.10 \mathrm{~m} \mathrm{CaSO}_{4}\) (b) the solubility of solids in water usually increases as the temperature increases. (c) pressure must be applied to cause reverse osmosis to occur. (d) \(0.10 \mathrm{M} \mathrm{BaCl}_{2}\) has a higher osmotic pressure than \(0.10 \mathrm{M}\) glucose. (e) molarity and molality are nearly the same in dilute solutions.

The freezing point of \(0.10 \mathrm{M} \mathrm{KHSO}_{3}\) is \(-0.38^{\circ} \mathrm{C}\). Which of the following equations best represents what happens when \(\mathrm{KHSO}_{3}\) dissolves in water? (a) \(\mathrm{KHSO}_{3}(s) \longrightarrow \mathrm{KHSO}_{3}(a q)\) (b) \(\mathrm{KHSO}_{3}(s) \longrightarrow \mathrm{K}^{+}(a q)+\mathrm{HSO}^{3-}(a q)\) (c) \(\mathrm{KHSO}_{3}(s) \longrightarrow \mathrm{K}^{+}(a q)+\mathrm{SO}_{3}{ }^{2-}(a q)+\mathrm{H}^{+}(a q)\)

Describe how you would prepare \(465 \mathrm{~mL}\) of \(0.3550 \mathrm{M}\) potassium dichromate solution starting with (a) solid potassium dichromate. (b) \(0.750 M\) potassium dichromate solution.

Show that the following relation is generally valid for all solutions: $$ \text { molality }=\frac{\text { molarity }}{d-\frac{\mathrm{MM}(\mathrm{molarity})}{1000}} $$ where \(d\) is solution density \(\left(\mathrm{g} / \mathrm{cm}^{3}\right)\) and \(\mathrm{MM}\) is the molar mass of the solute. Using this equation, explain why molality approaches molarity in dilute solution when water is the solvent, but not with other solvents.

An aqueous solution made up of \(32.47 \mathrm{~g}\) of iron(III) chloride in \(100.0 \mathrm{~mL}\) of solution has a density of \(1.249 \mathrm{~g} / \mathrm{mL}\) at \(25^{\circ} \mathrm{C}\). Calculate its (a) molarity. (b) molality. (c) osmotic pressure at \(25^{\circ} \mathrm{C}\) (assume \(\left.i=4\right)\). (d) freezing point.
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