Does vapour pressure of a liquid solution depend on the size of closed container, amount of solution taken, given that, temperature is kept constant? Or will it change, if some extra gas is added to the container at constant temperature?
The equilibrium vapor pressure of a compound at a given temperature is an intrinsic property of the compound, meaning it does not depend on variables like volume of the container, quantity of the compound, etc. Other examples of intrinsic properties include density and melting point. None of these intrinsic properties depends on the quantity of the material or it's physical surroundings.
The simple answer is that in a thermodynamically closed system composed of a condensed phase (either solid or liquid) and a vapor phase of a chemical compound (or element), then only temperature matters - at equilibrium. This means vapor pressure may not be the same as partial pressure - even at equilibrium - even if the temperature is the same. That is; despite Dalton's Law of Partial Pressures, the partial pressure of a compound does depend on the other components present in the gas phase. This effect is usually so slight, when dealing with pressures around STP, that it can be ignored for a good approximation and Dalton's Law can be assumed to hold. You ask if the amount of material matters. Well, only if it isn't enough to provide for any condensed phase to exist after equilibrium has been established. That is, yes a drop of water in a volume of a cubic meter at 150°C will only exist in the vapor phase, and it's pressure will not correspond to the vapor pressure of steam at 150°C. The partial pressure of a compound depends on temperature, gravitational gradient, surface curvature, and the interactions of the compound with the other materials present; but it is a good approximation to assume it depends only on temperature (assuming enough of the compound is present to have some of it in a condensed phase). There's also another complication, which is surface absorption (onto the walls of the container), but again, this can generally be ignored.
Vapor pressure is the pressure exerted by the vapors when in dynamic equilibrium with its condensed phase at a given temperature.
It is the maximum value of partial pressure at a given temperature. So if pressure is increased after this value then some vapors will condense and if pressure is decreased then more liquid will vaporise until the vapor pressure is reached.
Vapor pressure doesn't depend on the amount of the liquid, volume of vessel, shape/size of container, area of the surface, etc.
It only depends on the nature of the liquid and the temperature.
Let's compare vapor pressure of ethanol and ethylene glycol. Since intermolecular forces of attraction are less in ethanol as compared to glycol (More H-bonding in glycol), molecules of ethanol will find it easier to escape into gas phase and hence vapor pressure of ethanol is more than that of glycol. In short, vapor pressure is inversely proportional to forces of attraction.
Now let's see it's relation with temperature. As the temperature increases kinetic energy of the molecules increases and therefore vapor pressure increases. The relation is clearly described by the claussius clapeyron equation:
ln(P°) = -∆H/RT + c
Where P° is the vapor pressure
∆H is the enthalpy of vaporisation
R is universal gas constant
c is the constant of integration (check out the derivation)
d(lnP°)/dT = ∆H/RT²
By variable separable method of solving differential equations we obtain the above solution.
Using claussius clapeyron equation vapor pressure can be easily compared at two different temperatures.
Changing the composition will also change the vapor pressure. It depends on the nature of the solute. That can be given by Raoult's law for ideal solutions and the deviations in case of non ideal ones.
Hope this helps!