The solubility of N2 in blood at 37 °C and at a partial pressure of 0.80 atm is 5.6×10−4mol L−1. The Henry's law constant for this system is

5.6 × 10 − 4 atm L mol − 1

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The solubility of $N_{2}$ in blood at 37 °C and at a partial pressure of $0.80 a t m$ is $5.6 \times 10^{- 4} mol L^{- 1}$ . The Henry's law constant for this system is

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$1.43 \times 10^{3} atm L {mol}^{- 1}$

$2.86 \times 10^{2} atm L {mol}^{- 1}$

$\begin{aligned} 1.62 \times 10^{3} atm L {mol}^{- 1} \end{aligned}$

$5.6 \times 10^{- 4} atm L {mol}^{- 1}$

answer is A.

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Henry's law in terms of concentration of dissolved gas is $p = k_{H} c$ . Hence,

$k_{H} = \frac{p}{c} = \frac{0.80 atm}{5.6 \times 10^{- 4} mol L^{- 1}} = 1.43 \times 10^{3} atm L {mol}^{- 1}$

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