1. Homework Statement
If the boiling point is the point at which vapor pressure > atmospheric pressure, so all of the water molecules can break free and fly into the atmosphere (i.e. overcoming the atmospheric pressure), then why is it that when atmospheric pressure > vapor pressure, the atmospheric pressure is pushing down on the water such that no vapor can escape too?
I guess the question im asking is: Is vapor pressure independent of external pressure?
2. Homework Equations
3. The Attempt at a Solution
So Vapor pressure is the result of the phenomenon when the surface molecules of a liquid gain enough kinetic energy to leave the liquid and become a gas. The force per unit area of these molecules is the vapor pressure.
So it makes sense as T goes up, Vapor Pressure should increase since on average more molecules have the energy to break free from their intermolecular bonds.
From an intermolecular bond perspective this makes sense. But form a pressure perspective im still a bit confused. If the boiling point is the point at which the vapor pressure > the atmospheric pressure so all the molecules can keep leaving the surface, then how is it that molecules can leave the surface via evaporation when pAtmosphere > pVapor Pressure?
wouldnt the atmospheric pressure push all of the water molecules down into liquid? Just as all of the water molecules push against the atmospheric gas molecules to free themselves when pvapor pressure > patmosphere?
____________________________________
Also if im comparing the vapor pressure of a solid vs that of a liquid (sublimation vs evaporation), then if I decrease the temperature of both solid and liquid at the same rate BELOW the melting point of the solid (but at different pressures), apparently the vapor pressure of the liquid will decrease faster than that of the solid, to the extent where Vapor Pressure of the Solid > Vapor PRessure of the Liquid.
What's going on here? I assume its a different case because we're in two differnet phases?
If the boiling point is the point at which vapor pressure > atmospheric pressure, so all of the water molecules can break free and fly into the atmosphere (i.e. overcoming the atmospheric pressure), then why is it that when atmospheric pressure > vapor pressure, the atmospheric pressure is pushing down on the water such that no vapor can escape too?
I guess the question im asking is: Is vapor pressure independent of external pressure?
2. Homework Equations
3. The Attempt at a Solution
So Vapor pressure is the result of the phenomenon when the surface molecules of a liquid gain enough kinetic energy to leave the liquid and become a gas. The force per unit area of these molecules is the vapor pressure.
So it makes sense as T goes up, Vapor Pressure should increase since on average more molecules have the energy to break free from their intermolecular bonds.
From an intermolecular bond perspective this makes sense. But form a pressure perspective im still a bit confused. If the boiling point is the point at which the vapor pressure > the atmospheric pressure so all the molecules can keep leaving the surface, then how is it that molecules can leave the surface via evaporation when pAtmosphere > pVapor Pressure?
wouldnt the atmospheric pressure push all of the water molecules down into liquid? Just as all of the water molecules push against the atmospheric gas molecules to free themselves when pvapor pressure > patmosphere?
____________________________________
Also if im comparing the vapor pressure of a solid vs that of a liquid (sublimation vs evaporation), then if I decrease the temperature of both solid and liquid at the same rate BELOW the melting point of the solid (but at different pressures), apparently the vapor pressure of the liquid will decrease faster than that of the solid, to the extent where Vapor Pressure of the Solid > Vapor PRessure of the Liquid.
What's going on here? I assume its a different case because we're in two differnet phases?
At standard atmospheric pressure., water has gaseous form (water vapor) at 150°C Water freezes at 0 degrees Celsius. Below this temperature it is ice (solid). Water boils at 100 degrees Celsius.
Find the mole fraction of your solvent. The last thing we need to do before we can solve is to find the mole fraction of our solvent. Finding mole fractions is easy: just convert your components to moles, then find what percentage of the total number of moles in the substance each component occupies. In other words, each component's mole fraction equals (moles of component)/(total number of moles in the substance.)
- Let's say that our recipe for simple syrup uses 1 liter (L) of water and 1 liter of sucrose (sugar.) In this case, we'll need to find the number of moles in each. To do this, we'll find the mass of each, then use the substance's molar masses to convert to moles.
- Mass (1 L of water): 1,000 grams (g)
- Mass (1 L of raw sugar): Approx. 1,056.7 g[7]
- Moles (water): 1,000 grams × 1 mol/18.015 g = 55.51 moles
- Moles (sucrose): 1,056.7 grams × 1 mol/342.2965 g = 3.08 moles (note that you can find sucrose's molar mass from its chemical formula, C12H22O11.)
- Total moles: 55.51 + 3.08 = 58.59 moles
- Mole fraction of water: 55.51/58.59 = 0.947