How size of molecules and intrmolecular forces affect vapour pressure?
927912
Bigger molecule means stronger intermolecular forces, and as a resault lower vapour pressure at a given temperature.
523715
Higher Velocity will cause lower pressure and larger volume that will decrease its temperature. Volume, temperature and pressure are all physically related in gas.
811149
Gases are ultimately soluble in each other as noted by the atmosphere. If the pressure is increased, the molecules will get closer together.  But, the solubility stays the same.
227415
actual pressure of the vapour\ndivided by\nthe maximum pressure of the vapour at 'that' temperature\nequals (in this case) 1 = 100%
950681
Van derWaal's equation is this:\n(P+a(n/V)^2)(V-nb)=nRT\n\nThe ideal gas law is:\nPV=nRT\n\na) b will lower the total volume, accounting for the space taken up by the molecules.  True.\n\nb) a will increase the volume based on intermolecular repulsive forces.  True.\n\nc) The attraction will decrease the pressure.  False.\n\nd) The individual particles will stay a constant size; the volume and pressure don't change bond lengths or shell sizes, just the spacing between the particles.  True.
499971
When gas is compressed in a thermally insulated vessel the gas is heated.\n\nWhen pressure increases temperature increases.\n\nSimilarly if the gas is expanded in a thermally insulated vessel the gas is cooled.\n\nWhen the pressure decreases the temperature decreases. \n\nNear earth’s surface, the atmospheric pressure is very high.\n\nAt a height of 30km from the surface, the pressure of air is almost zero.\n\nThough the hot air goes up; as it goes up the pressure decreases and volume increases (density decreases) and as per the  gas equation the temperature decreases.
316023
intermolecular polarity and the fact that fluids dont compress the way gases do?
149015
PV=nRT is the formula you want to use to explain this. P is Pressure, V = Volume, n= Moles of molecule, R is a constant, T = Temerature. Now the reasoning behind this is because molecules that in the in gas phase want to be far apart from each other as possible. If you try to increase the pressure, you are forcing the molecules to come closer together. These molecules then rub up against each other more, creating heat.
112298
Ideal gas laws assume that all particles are spherically symmtric and engage only in elastic collisions.  They ignore all other types of molecular energy (rotation in other planes of symmetry or asymmetry, bond stretching/compressing/bending, rotation about a molecular bond--all these things can and do absord energy from a collision), and they negate the effects of all intermolecular forces, like dipole interations, hydrogen bonding, London dispersion forces, etc.\n\nIn gases like helium, the atoms are basically spherically symmetric, and most of these random forces are neglible.  Hence, ideal gas laws work really well for these--only at really ridiculously low temperatures and high pressures do these forces manage to liquify helium.  The larger you molecule is--and I'm not talking about mass so much as the number of chemical bonds--the less ideally it behaves.  The greater the heterogeneity of a gass, the less ideally it behaves.  The greater the polarity (eg seperation of charges) displayed by the molecule, the less ideally it behaves.
1156050
I would visit the site cited below.  The discussion includes:\n\n"Irving Langmuir developed a way to measure vapor pressure \nby measuring the evaporation rate, which we will now flip around backwards. His reasoning is that the rate at which molecues are lost due to evaporation to a gas with no partial pressure of the evaporating substance is the same as the rate at which molecules of the substance would hit the surface if it were in equilibrium with the vapor (because in equilibrium the evaporation rate and re-condensation rate cancel each other out). His expression is: \n\n(mass loss rate)/(unit area) = (vapor pressure - ambient partial pressure)*sqrt( (molecular weight)/(2*pi*R*T) ) "
606777
the actual equation is:  (p + b·n^2/V^2)(V - a·n) = nRT\n\na:- every mol of particles of the gas has a tiny finite volume of a\n\nb:-modification is due to the particles sticking to each other, a phenomenon called cohesion\n\nR- rydberg's constant
127151
Hot air has molecules which are more active from the energy that they receive from heat and, therefore, tend to collide more often with other molecules than when the air is cool. When more collisions occur, more pressure will be built up. To maintain pressure in the surroundings, the molecules will move to a place with less pressure, and that is upwards. When the air cools, it will go down again and the cycle continues.
450793
It is because the varius atoms and molecules have different sizes, shapes and weights. This means it take different amount of energy to make them vaporize.
336395
I'm not pretty sure about this but because of the low molecular weight, the attraction between molecules is little and therefore outside forces can easily scatter these molecules troughout space. This results in gas state rather than liquid or solid
141535
The Climate.Pressure increases if temperature increases and volume stays the same. This is because the molecules are moving faster and colliding harder with the sides of the container. The greater the mass of air in the container, the greater the number of molecules in the container and the higher its density. The air molecules in the container collide against the sides of the container, so the more of them there are colliding against sides of the container the greater the overall force on the sides of the container. If the temperature is high enough the moisture is vapor or steam and if the moisture doesn't condense on the sides of the container then the moisture molecules are also colliding with the side of the container. On the other hand, if the air in the container is cool enough that the moisture does condense then the moisture is no longer exerting outward force on the container (or at least there's much less of it still moving around as vapor and colliding with the sides of the container).\n\nWhen you're talking about a body of water with gas (carbon dioxide) and sugar in it, like Coke, if the bottle is still closed then you have a volume of air above the soda that exerts enough force on the soda to prevent the gas from coming out of the soda at all until you open the bottle. Once the bottle is open, think about surface tension. That's the reason the carbon dioxide is released slowly when you open the bottle. If you disrupt the surface tension the gas can release all at once.
433864
Use the General Gas Equation:\n\nPV = nRT\n\nwhere\nP = pressure \nV = volume \nn = number of moles\nR = Universal Gas Constant (value depends on units of P,V, n, T eg R = 62.3638 when V is in litres, P is in mmHg, T in Kelvins and n in moles.)\n\nUse the above equation to solve for n. If you want the answer in grams, simply multiply the number of moles by the molar mass of H2.\n\nIf you require a more precise answer, then you should use the Van der Waals equation (a more sophisticated version of the General Gas Equation that takes into account interactions between particles and the volume of those particles).
879801
Pentane\n\nIt depends on which compound has the strongest intermolecular forces. In the case of hydrocarbons the intermolecular forces that are present are dispersion forces.\nIn such cases:\n"The more branched a molecule is, the more "spherical" it is, and the smaller surface area it will have - since dispersion forces increase with surface area, the more branched molecules will have a lower boiling point. You may notice that the trend in the melting points is rather different - this is because the longer a branch is, the more potential it has to flop around and disrupt the regular crystalline lattice of frozen pentane, giving a liquid state."