question: Why does water evaporate even if ambience temp is less than its boiling point?
comment: ELI5 - think of liquid water molecules like birds standing on the ground at a park. to get all of the birds to get scared and fly into the air (become water vapor), a loud enough sound needs to be made (temperature). If that loud sound gets made, they all fly away (evaporate). However, even if a loud sound isn't made, some of the birds are going to get restless and fly off on their own. Eventually, even if a loud sound isn't made, enough birds will get restless and fly away that none will be left More sciencey - whether H20 is in the liquid state (water) or gas state (vapor) depends on the energy of the H20 molecules. Heating them up imparts energy onto them, and they shake more and get spread apart and evaporate as a gas. However, individual H20 water molecules can randomly escape as a gas without heat. Basically, if you have a bowl of water, individual water molecules at the top are escaping/evaporating into gas. Eventually, they all evaporate. Try to think of it as a constant process. At the top of the bowl of water, some molecules are changing into gas while some are changing into liquid. At room temp, the process favors gas formation, so it eventually all becomes gas.
passages: On a molecular level, there is no strict boundary between the liquid state and the vapor state. Instead, there is a Knudsen layer, where the phase is undetermined. Because this layer is only a few molecules thick, at a macroscopic scale a clear phase transition interface cannot be seen. 
   In the "vicinity" of the critical point, the physical properties of the liquid and the vapor change dramatically, with both phases becoming ever more similar. For instance, liquid water under normal conditions is nearly incompressible, has a low thermal expansion coefficient, has a high dielectric constant, and is an excellent solvent for electrolytes. Near the critical point, all these properties change into the exact opposite: water becomes compressible, expandable, a poor dielectric, a bad solvent for electrolytes, and prefers to mix with nonpolar gases and organic molecules.
   If evaporation takes place in an enclosed area, the escaping molecules accumulate as a vapor above the liquid. Many of the molecules return to the liquid, with returning molecules becoming more frequent as the density and pressure of the vapor increases. When the process of escape and return reaches an equilibrium, the vapor is said to be "saturated", and no further change in either vapor pressure and density or liquid temperature will occur. For a system consisting of vapor and liquid of a pure substance, this equilibrium state is directly related to the vapor pressure of the substance, as given by the Clausius–Clapeyron relation:
   These results led to the problem of how new phases can ever arise from old ones. For example, if a container filled with water vapour at slightly below the saturation pressure is suddenly cooled, perhaps by adiabatic expansion, as in a cloud chamber, the vapour may become supersaturated with respect to liquid water. It is then in a metastable state, and we may expect condensation to take place. A reasonable molecular model of condensation would seem to be that two or three molecules of water vapour come together to form a tiny droplet, and that this nucleus of condensation then grows by accretion, as additional vapour molecules happen to hit it. The Kelvin equation, however, indicates that a tiny droplet like this nucleus, being only a few ångströms in diameter, would have a vapour pressure many times that of the bulk liquid. As far as tiny nuclei are concerned, the vapour would not be supersaturated at all. Such nuclei should immediately re-evaporate, and the emergence of a new phase at the equilibrium pressure, or even moderately above it should be impossible. Hence, the over-saturation must be several times higher than the normal saturation value for spontaneous nucleation to occur.
   where "P", "P" are the vapor pressures at temperatures "T", "T" respectively, Δ"H" is the enthalpy of vaporization, and "R" is the universal gas constant. The rate of evaporation in an open system is related to the vapor pressure found in a closed system. If a liquid is heated, when the vapor pressure reaches the ambient pressure the liquid will boil.
   Under certain conditions, such as when the boiling temperature of water is reached, a net evaporation will always occur during standard atmospheric conditions regardless of the percent of relative humidity. This immediate process will dispel massive amounts of water vapor into a cooler atmosphere.
   The amount of water vapor in an atmosphere is constrained by the restrictions of partial pressures and temperature. Dew point temperature and relative humidity act as guidelines for the process of water vapor in the water cycle. Energy input, such as sunlight, can trigger more evaporation on an ocean surface or more sublimation on a chunk of ice on top of a mountain. The "balance" between condensation and evaporation gives the quantity called vapor partial pressure.
   These two definitions are equivalent: the boiling point is the temperature at which the increased entropy of the gas phase overcomes the intermolecular forces. As a given quantity of matter always has a higher entropy in the gas phase than in a condensed phase (formula_6 is always positive), and from
  
  the Gibbs free energy change falls with increasing temperature: gases are favored at higher temperatures, as is observed in practice.
  
  Section::::Vaporization enthalpy of electrolyte solutions.
   Evaporation is a phase transition from the liquid phase to vapor (a state of substance below critical temperature) that occurs at temperatures below the boiling temperature at a given pressure. Evaporation occurs "on the surface". Evaporation only occurs when the partial pressure of vapor of a substance is less than the equilibrium vapor pressure. For example, due to constantly decreasing or negative pressures, vapor pumped out of a solution will leave behind a cryogenic liquid.
   Data in the table above is given for water–steam equilibria at various temperatures over the entire temperature range at which liquid water can exist. Pressure of the equilibrium is given in the second column in kPa. The third column is the heat content of each gram of the liquid phase relative to water at 0 °C. The fourth column is the heat of vaporization of each gram of liquid that changes to vapor. The fifth column is the work "P"Δ"V" done by each gram of liquid that changes to vapor. The sixth column is the density of the vapor.
   At room temperature and pressure, the water jar reaches equilibrium when the air over the water has a humidity of about 3%. This percentage increases as the temperature goes up. At 100 °C and atmospheric pressure, equilibrium is not reached until the air is 100% water. If the liquid is heated a little over 100 °C, the transition from liquid to gas will occur not only at the surface, but throughout the liquid volume: the water boils.
  
  Section::::Number of phases.
   Since the saturation vapor pressure is proportional to temperature, cold air has a lower saturation point than warm air. The difference between these values is the basis for the formation of clouds. When saturated air cools, it can no longer contain the same amount of water vapor. If the conditions are right, the excess water will condense out of the air until the lower saturation point is reached. Another possibility is that the water stays in vapor form, even though it is beyond the saturation point, resulting in supersaturation.
   If the vapor then condenses to a liquid on a surface, then the vapor's latent energy absorbed during evaporation is released as the liquid's sensible heat onto the surface.
  
  The large value of the enthalpy of condensation of water vapor is the reason that steam is a far more effective heating medium than boiling water, and is more hazardous.
  
  Section::::Usage.:Meteorology.
   Viewed from a microscopic perspective, the dissolution and dilution processes involve three steps of molecular interaction: the breaking of attraction between solute molecules (lattice energy), the breaking of attraction between solvent molecules, and the forming of attraction between a solute and a solvent molecule. If the solution is ideal, which means the solute and the solvent are identical in an interaction, then all the kinds of attraction mentioned above have the same value. As a result, the enthalpy change caused by breaking and forming attraction is canceled, and the dilution of an ideal solution causes no enthalpy change.
   Gaseous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water (to minimum, which is below 120 °C for most permanent gases), but more soluble in organic solvents.
   Section::::Properties.:Respiration and breathing.
  
  Water vapor is a by-product of respiration in plants and animals. Its contribution to the pressure, increases as its concentration increases. Its partial pressure contribution to air pressure increases, lowering the partial pressure contribution of the other atmospheric gases (Dalton's Law). The total air pressure must remain constant. The presence of water vapor in the air naturally dilutes or displaces the other air components as its concentration increases.
   where "z" is the height of the protein clump in the bottle and "g" is the acceleration due to gravity. Since "s" = 1, the average potential energy of a protein clump equals "k""T". Hence, a protein clump with a buoyant mass of 10 MDa (roughly the size of a virus) would produce a haze with an average height of about 2 cm at equilibrium. The process of such sedimentation to equilibrium is described by the Mason–Weaver equation.
  
  Section::::History.
   Section::::Table of specific latent heats.
  
  The following table shows the specific latent heats and change of phase temperatures (at standard pressure) of some common fluids and gases.
  
  Section::::Specific latent heat for condensation of water in clouds.
  
  The specific latent heat of condensation of water in the temperature range from −25 °C to 40 °C is approximated by the following empirical cubic function:
  
  where the temperature formula_4 is taken to be the numerical value in °C.
   After relaxation, the speed of some low mass members can be greater than the escape velocity of the cluster, which results in these members being lost to the cluster. This process is called evaporation. (A similar phenomenon explains the loss of lighter gases from a planet, such as hydrogen and helium from the Earth—after equipartition, some molecules of sufficiently light gases at the top of the atmosphere will exceed the escape velocity of the planet and be lost.)
   Condensation is initiated by the formation of atomic/molecular clusters of that species within its gaseous volume—like rain drop or snow flake formation within clouds—or at the contact between such gaseous phase and a liquid or solid surface. In clouds, this can be catalyzed by water-nucleating proteins, produced by atmospheric microbes, which are capable of binding gaseous or liquid water molecules.
  
  Section::::Reversibility scenarios.
  
  A few distinct reversibility scenarios emerge here with respect to the nature of the surface.
  
  BULLET::::- absorption into the surface of a liquid (either of the same substance or one of its solvents)—is reversible as evaporation.
   For example, water has a critical temperature of , which is the highest temperature at which liquid water can exist. In the atmosphere at ordinary temperatures, therefore, gaseous water (known as water vapor) will condense into a liquid if its partial pressure is increased sufficiently.
  
  A vapor may co-exist with a liquid (or a solid). When this is true, the two phases will be in equilibrium, and the gas-partial pressure will be equal to the equilibrium vapor pressure of the liquid (or solid).
  
  Section::::Properties.
   For example, if the vapor's partial pressure is 2% of atmospheric pressure and the air is cooled from 25 °C, starting at about 22 °C water will start to condense, defining the dew point, and creating fog or dew. The reverse process accounts for the fog burning off in the morning. If the humidity is increased at room temperature, for example, by running a hot shower or a bath, and the temperature stays about the same, the vapor soon reaches the pressure for phase change, and then condenses out as minute water droplets, commonly referred to as steam.
   Separation of components ("e.g." water and ethanol) is based on a difference in transport rate of individual components through the membrane. This transport mechanism can be described using the solution-diffusion model, based on the rate/ degree of dissolution of a component into the membrane and its velocity of transport (expressed in terms of diffusivity) through the membrane, which will be different for each component and membrane type leading to separation.
  
  Section::::Applications.
   Evaporative cooling (atomic physics)
  
  Evaporative cooling is an atomic physics technique to achieve high phase space densities which optical cooling techniques typically can not reach.
   Water in a closed jar with an air space over it forms a two phase system. Most of the water is in the liquid phase, where it is held by the mutual attraction of water molecules. Even at equilibrium molecules are constantly in motion and, once in a while, a molecule in the liquid phase gains enough kinetic energy to break away from the liquid phase and enter the gas phase. Likewise, every once in a while a vapor molecule collides with the liquid surface and condenses into the liquid. At equilibrium, evaporation and condensation processes exactly balance and there is no net change in the volume of either phase.
  
presuppositions: Water needs to be boiling to evaporate. Water needs to be boiling to evaporate.
corrections: Water at room temp is constantly changing into a gas. Water at room temp is constantly changing into a gas.
raw_presuppositions: Water needs to be boiling to evaporate. Water needs to be boiling to evaporate. Water needs to be boiling to evaporate.
raw_corrections: Water at room temp is constantly changing into a gas. Water at room temp is constantly changing into a gas. Water at room temp is constantly changing into a gas.