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27 Nov 2019

The Henry’s law constant for O2 is 1.3 × 10−3 M/atm at 25 °C. What mass of oxygen would be dissolved in a 40-L aquarium at 25 °C, assuming an atmospheric pressure of 1.00 atm, and that the partial pressure of O2 is 0.21 atm?

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Hubert Koch
Hubert KochLv2
9 Apr 2019

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Related questions

The Henry’s law constant for O2 is 1.3×10^−3 M/atm at 25 °C. What mass of oxygen would be dissolved in a 40-L aquarium at 25 °C, assuming an atmospheric pressure of 1.00 atm, and that the partial pressure of O2 is 0.21 atm?

Given that the solubility of O2 in water at 25oC is 8.70 ppm, calculate (a) the concentration of dissolved O2 in mol L-1 and (b) the value of the Henry’s law constant, in mol L-1 kPa-1, for O2 at this temperature? The atmospheric pressure is 1.00 atm. (Hint: the mole fraction of oxygen in air is 0.210).

When pure water is exposed to atmospheric CO2, some of the gas dissolves in the water to create H2CO3. The partial pressure of CO2 in the atmosphere is approximately 4.0E-4 atm. The Henry’s law constant (k) for CO2 at 25°C is 3.2E-2 mol/(L*atm). What is the pH of pure water exposed to atmospheric CO2? Assume that pH is controlled by the first proton dissociation of H2CO3.

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