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Chapter 11: Problem 50

Calculate the solubility of \(\mathrm{O}_{2}\) in water a partial pressure of \(\mathrm{O}_{2}\) of 120 torr at \(25^{\circ} \mathrm{C}\). The Henry's law constant for \(\mathrm{O}_{2}\) is \(1.3 \times 10^{-3} \mathrm{~mol} / \mathrm{L} \cdot\) atm for Henry's law in the form \(C=k P\), where \(C\) is the gas concentration \((\mathrm{mol} / \mathrm{L})\).

Short Answer

Expert verified
The solubility of O₂ in water at a partial pressure of 120 torr and \(25^{\circ}C\) is approximately \(2.05 \times 10^{-4}\: \mathrm{mol} / \mathrm{L}\).

Step by step solution

01

Pressure conversion formula

\(P_{atm} = \frac{P_{torr}}{760}\) #Step 2: Calculate pressure in atm# Now we'll plug in the given pressure value and calculate the pressure in atm.
02

Pressure in atm calculation

\(P_{atm} = \frac{120 \: torr}{760 \: torr} = 0.1579 \: atm\) #Step 3: Use Henry's law to find solubility# Now that we have the pressure value in atm, we'll use Henry's law formula to find the solubility of O₂ in water.
03

Henry's law formula

\(C = kP\) Here, the Henry's law constant, k, is given as \(1.3 \times 10^{-3}\: \mathrm{mol} / \mathrm{L} \cdot \mathrm{atm}\) #Step 4: Calculate solubility of O₂ in water at given pressure# Finally, we'll use the values of k and P to find the solubility of O₂ in water (C).
04

Solubility calculation

\(C = (1.3 \times 10^{-3}\: \mathrm{mol} / \mathrm{L} \cdot \mathrm{atm}) \times 0.1579 \: atm\) \(C = 2.05 \times 10^{-4}\: \mathrm{mol} / \mathrm{L}\) Hence, the solubility of O₂ in water at a partial pressure of 120 torr and \(25^{\circ}C\) is approximately \(2.05 \times 10^{-4}\: \mathrm{mol} / \mathrm{L}\).

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Most popular questions from this chapter

From the following: pure water solution of \(\mathrm{C}_{12} \mathrm{H}_{22} \mathrm{O}_{11}(m=0.01)\) in water solution of \(\mathrm{NaCl}(m=0.01)\) in water solution of \(\mathrm{CaCl}_{2}(m=0.01)\) in water Choose the one with the a. highest freezing point. b. lowest freezing point. c. highest boiling point. d. lowest boiling point. e. highest osmotic pressure.

An aqueous solution is \(1.00 \% \mathrm{NaCl}\) by mass and has a density of \(1.071 \mathrm{~g} / \mathrm{cm}^{3}\) at \(25^{\circ} \mathrm{C}\). The observed osmotic pressure of this solution is \(7.83\) atm at \(25^{\circ} \mathrm{C}\).

In lab you need to prepare at least \(100 \mathrm{~mL}\) of each of the following solutions. Explain how you would proceed using the given information. a. \(2.0 \mathrm{~m} \mathrm{KCl}\) in water (density of \(\mathrm{H}_{2} \mathrm{O}=1.00 \mathrm{~g} / \mathrm{cm}^{3}\) ) b. \(15 \% \mathrm{NaOH}\) by mass in water \(\left(d=1.00 \mathrm{~g} / \mathrm{cm}^{3}\right)\) c. \(25 \% \mathrm{NaOH}\) by mass in \(\mathrm{CH}_{3} \mathrm{OH}\left(d=0.79 \mathrm{~g} / \mathrm{cm}^{3}\right)\) d. 0. 10 mole fraction of \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}\) in water \(\left(d=1.00 \mathrm{~g} / \mathrm{cm}^{3}\right)\)

Calculate the sodium ion concentration when \(70.0 \mathrm{~mL}\) of 3.0 \(M\) sodium carbonate is added to \(30.0 \mathrm{~mL}\) of \(1.0 \mathrm{M}\) sodium bicarbonate.

The high melting points of ionic solids indicate that a lot of energy must be supplied to separate the ions from one another. How is it possible that the ions can separate from one another when soluble ionic compounds are dissolved in water, often with essentially no temperature change?
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