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Partial pressures are basically the pressure of each gas if it were to occupy the container just by itself (the total pressure of a container of mixed gases is the sum of all the partial pressures of each gas). Since concentration is moles/liters, you can technically put moles of gas over liters. Hope this helps!
Partial pressure is like the force exerted by a gas if it was just by itself in some volume. In PV=nRT, we can manipulate it to be: P/RT = n/v. n/V is concentration, and we can see that Pressure is proportional to concentration. 1/RT is our proportionality constant.
Brian_Wu_3B wrote:What do partial pressures really mean? Why can they be converted to concentrations? I'm just not quite understanding what the term means.
The pressure of a mixture of gases is the sum of the partial pressures of the different gases in the mixture. So we're basically saying that all the gases in the mixture contribute a portion to the total pressure, or a partial pressure. When calculating your equilibrium constant either concentrations or partial pressures can be utilized depending on the components of the mixture so there is a relationship between Kp and Kc : Kp =Kc (RT)^(delta n)
delta n is the change in moles of gases.
Partial pressures are the molar fraction of a gas over total pressure in a container (from adding up the total moles of gas in the reaction). I believe in all of the examples concerning equilibrium, the total pressure is 1 atm.
Each gas in an equilibrium reaction takes up space/volume, meaning that they each exert their own partial pressure by colliding with other particles and their container. The sum of all of these partial pressures results in the total pressure, which is the amount of pressure exerted on the container by all of the gases inside of it. We can convert partial pressures to concentrations because molar concentrations are measures of moles/liter, so we can calculate the concentrations of each gas in a reaction given their partial pressure by rearranging PV=nRT to n/V=P/RT. Hope this helps
One way to think of it is by comparing it to concentrations. Similar to how each element has its own unique concentration in a solution. Each element has its own unique partial pressure. Hope this helps!
I understand why we can convert between partial pressure and concentration but it seemed to me that prof Lavelle used them interchangeably in lecture. Can they be used interchangeably or does it just depend on the context and I just wasn't understanding the context he was using it in?
Partial pressures is the pressure of each individual gas. When all the partial pressures are added up, you get the total pressure. It can be convereted to concentration (which is measured in mol/L) using the ideal gas law: PV = nRT --> you solve to get n/V which would be mol/L.
Partial pressure is just the pressure the gas would have in the container if it was by itself. All the partial pressures add up to the total pressure of the container. But each gas is exerting its own pressure on the container. Hope that helps!
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