Background Nitrogen, the most abundant gas in the atmosphere (78%), is not necessary in the cardiovascular system and it causes no harm under typical atmospheric conditions. For instance, at standard atmospheric pressure nitrogen can absorb into your skin. This is not an issue until you dive underwater. Henry's Law (chemistry) states that the amount of gas that dissolves in a liquid is directly proportional to the gas' partial pressure. When the partial pressure of nitrogen increases, it will be absorbed into the blood stream. This is not particularly dangerous, although it can lead to a pleasant feeling called "nitrogen narcosis" in some people. The real danger is when the diver resurfaces. When the diver begins to ascend, the partial pressure of nitrogen decreases and nitrogen comes out of solution as small bubbles of gas in the bloodstream. The buildup of nitrogen can be removed from the body by passing nitrogen through the bloodstream and into the lungs where it is exhaled. However, this takes time. If a diver surfaces quickly, problems can arise. (a) A recreational diver descends to 220 feet and stays there for some time. What absolute pressure do they experience at this depth? Tip: Make sure you convert feet to m! P Pa depth (b) As the diver begins to ascend, a small bubble of nitrogen (2.1 μm diameter) comes out of solution into the bloodstream. What is the diameter of this bubble at the surface? Hint: Use the Ideal Gas Law assuming constant temperature (P;V; = P+V+). What is the volume of a sphere? Diameter = μη (c) How does this diameter compare to the diameter of the capillaries (6 - 8 µm) or venules (10-50 μm) of the body? Will this bubble cause problems? (d) Usually the bubbles don't stay separated. If many bubbles of nitrogen come out of solution, they will cluster together. If the bubbles cluster such that a 'super-bubble' of 30 μm diameter forms at a depth of 220 feet, will this cause problems?

Learning Goal: To understand the ideal gas law and be able to apply it to a wide variety of situations. The absolute temperature T, volume V, and pressure p of a gas sample are related by the ideal gas law, which states that PV = nRT Here n is the number of moles in the gas sample and R is a gas constant that applies to all gases. This empirical law describes gases well only if they are sufficiently dilute and at a sufficiently high temperature that they are not on the verge of condensing. In applying the ideal gas law, p must be the absolute pressure, measured with respect to vacuum and not with respect to atmospheric pressure, and I must be the absolute temperature, measured in kelvins (that is, with respect to absolute zero, defined throughout this tutorial as -273°C). If p is in pascals and V is in cubic meters, use R = 8.3145 J/(mol · K). If p is in atmospheres and V is in liters, use R = 0.08206 L atm/(mol-K) instead. Part A A gas sample enclosed in a rigid metal container at…

Can you please solve the below problem, showing step by step: A balloon contains mostly helium and a little nitrogen.  The pressure of the helium is 425 mm Hg.  What is the partial pressure of the nitrogen if the total pressure of the balloon is 67.98 kPa?