In order to be a gas does the pressure exerted by the gas has to be more than the atmospheric pressure?

In order to be a gas does the pressure exerted by the gas has to be more than the atmospheric pressure?  

Really simple, I think yes to this question but teacher says no. So what is your view into this?




Your teacher is right. For example, at room temperature, the pressure of water vapor over a pond or lake might be around 25 torr; that's less than atmospheric pressure, but the water vapor is undeniably a gas.


Maybe your confusion stems from a belief that gases must be produced by boiling liquids (since a liquid must have a vapor pressure equal to the external pressure before it begins to boil). But a liquid can also become a gas through evaporation, no matter what the external pressure is.


In **evaporation**, molecules escape from the liquid's surface and become a gas.  

Molecules in the liquid are in constant motion, and they frequently collide with each other. Every now and then one will get knocked off the surface of the liquid by the others. This can happen at any temperature or pressure the liquid exists at.


In **boiling**, bubbles of gas form throughout the liquid. This can only happen when the temperature is high enough to make the vapor pressure equal to the external pressure.


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No. Suppose we had a container containing nothing but water. Suppose it is only half full, so there is empty space. **There is no pressure from other gases.** If we heat this container enough, the water will boil and eventually turn completely into a gas. 
In short, clearly other things influence the state of matter in everyday life besides pressure from atmospheric gases. In this example, I illustrate how water evaporating in a vacuum is determined by temperature. I think this is a good example because it involves just one substance and doesn't bring into question how to think about mixed gases and partial pressures. 


There is an extremely useful equation that connects these relationships beautifully: 


$PV = nRT$


P stands for pressure, V for volume, T for temperature, n for the number of molecules, and R is a constant. 


This equation basically says that
*any change in pressure or volume to a closed system of gas will result in a corresponding change to the temperature of the system.* 


Thus, as long as $P\_1V\_1\neq P\_2V\_2$, you can always think about a change in volume or pressure as a change in temperature for a closed system. 


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If you know Dalton's law of partial pressure that "in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases"
And since pressure is non-negative quantity then:


P(total)=P1 + P2 + ... + Pn


SO the pressure of gas in atmosphere in fact, cannot be greater than atmospheric pressures . 


<http://en.wikipedia.org/wiki/File:Dalton%27s_law_of_partial_pressures.png>


And also I am not saying that gas stored in container cannot have pressure greater than atmosphere . The best example is your LPG cylinder.


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