How to find the number of moles in a gas,contact number to order sky,how to make my phone ring longer 02 - Tips For You

We have a pretty good understanding of the gaseous state in terms of the link between the variables describing the microscopic (atomic-scale) realm, and the variables describing the macroscopic realm. In the microscopic world we can talk about position, velocity, and mass of the individual atoms, all of which are governed by Newton's laws of motion. We start with the experimental observations of Robert Boyle, who showed that the quantities of pressure times volume are a constant for a sample at equilibrium. For an ideal gas the pressure times the volume is constant, but for real-world gases this is not always true. The straight horizontal line in the $p V$ versus $p$ graph is for an ideal gas, while the other two lines are those of some real gases.
Okay, then what causes the pV produce to be smaller than the ideal value as the pressure initially increases? And what causes the pV product to be larger than the ideal value as the pressure gets much higher? Therefore, when you hold your breath, you create a closed system with your lungs and thus Boyle's law will hold. In the plot of $V$-vs-$T$ above, we observe that the relationship is linear, with an intercept of absolute zero on the Kelvin scale.

Emile Clapeyron put Boyle's, Charles', and Avogadro's laws together and came up with the equation of state that we discussed earlier. We can also use the ideal gas law as a way of converting volumes into moles when calculating stoichiometric quantities in chemical reactions. To calculate the mass of a higher number of moles, or even calculate the mass of a certain number of moles, a formula triangle can be used. The formula mass of calcium carbonate can be calculated using the formula and values from the data book. By the use of statistical mechanics and averages we can then relate these variables to the macroscopic world by talking about such quantities as pressure, volume, temperature, and moles.
The thing to notice here is that as pressure goes to zero all gases become ideal, and this is shown by the fact that as $p$ goes to zero, all lines converge to the same point.
Something that has a high density, like a solid, has a very low compressibility since it takes a lot of pressure to change the volume even slightly, whereas something that has a low density, such as a gas, has a very high compressibility.
It is well known among divers that diving at the surface can be more dangerous than deeper diving. If you are down at 90 meters (at 10 atm) and you rise 10 meters to 80 meters (at 9 atm), the pressure has decreased by about 10%, and since $pV$ is a constant your lungs expand by about 10% (probably not too bad).

For now, though, let's focus on the macroscopic world and search for an equation of state that will tell us the relationships between all these macroscopic world quantities. We can understand this by first noting that for every 10 meters you descend in the water, the pressure increases by about 1 atm. Now if you are at 10 meters (at 2 atm), and you rise 10 meters to the surface (at 1 atm) the pressure had decreased by 50% and expanding your lungs by this factor could cause significant damage, maybe death!
The dashed lines on the graph are to represent the fact that you can not achieve absolute zero.
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