UCSB Science Line
Sponge Spicules Nerve Cells Galaxy Abalone Shell Nickel Succinate X-ray Lens Lupine
UCSB Science Line
How it Works
Ask a Question
Search Topics
Our Scientists
Science Links
Contact Information
Why does the water condense after evaporating?
Question Date: 2014-11-04
Answer 1:

It is possible for water to condense after evaporating if the correct conditions are met. For instance, let us consider how clouds form.

The few degrees of latitude directly above and below the equator receive the most direct sunlight throughout the year. Because of this, the temperatures over the ocean there get quite high; high enough for water to evaporate from the ocean. Water molecules rise in the form of gas from the surface of the ocean into the sky, and as they rise higher and higher, the temperature of the environment decreases. Eventually, the temperature is cool enough that the gas molecules condense and form little liquid droplets, agglomerated with dust, sand, and other airborne particles. These water molecules will undergo many more cycles of evaporation/condensation as air currents move them from the equatorial region towards the poles.

Answer 2:

Water evaporates when is gets enough energy in the form of heat to go from liquid to vapor. At room temperature, there is a small portion of water molecules that have enough energy to evaporate whereas in a boiling liquid, many more molecules have enough energy. To stay in the vapor phase, the water must stay warm enough. However, when the water vapor collides with a cold object, it can immediately get cooled by this object which will cause it to lose enough energy to become water again. This is why a cold water bottle will often have water on the outside because any water vapor in the air that bumps in to the water bottle will turn back into liquid water on the outside of the bottle.

Answer 3:

Thank you for asking such a great question. You want to know why water condenses after it has already evaporated. First, why does water evaporate? The water molecules at the surface of a liquid are moving fairly quickly and sometimes they have enough motion that they get knocked into the air. This event is evaporation. Only a specific amount of water molecules can be in the gaseous state at any time.

The amount of water molecules in vapor form is dependent on the temperature. Colder temperatures allow for less water vapor than warmer temperatures. Second, why does water condense? Water will condense when the temperature drop suddenly and the amount of water than can exist in the gaseous state drops. The water vapor turns into liquid water. For example, if you have a cold beverage on a hot day and humid day. The hot temperatures allow for a large amount of water vapor, but near the surface of the cold glass the temperature drops. The temperature is low enough for a large amount of water vapor to no longer be in the gaseous state. The water condenses on the outside of the glass.

To answer your question, if water evaporates but then enters an area which is too cold to keep it in the gaseous state, it will condense.

Answer 4:

Water vapor in the air condenses into liquid at a rate that depends on both temperature and the density of water vapor in the air; the warmer the air, the slower the condensation. Water evaporates off of a liquid, also at a rate that depends on temperature. Both processes happen at the same time. If the temperature is warm enough that the rate of evaporation is faster than the rate of condensation, then any water will slowly evaporate. If the rate of condensation is faster, however, then it will condense.

Normally, when you see condensation, the water was evaporated off of a warm liquid and is condensing onto the surface of a cold object, such as the outside of a glass if ice-water or something like it. The cold temperature on the surface of the glass is what prevents the water on the glass from evaporating off of it, which is why it stays in a liquid form on the surface. Water also condenses to form clouds in the sky because the air is colder at higher altitude.

Answer 5:

Thank you for the question! Air can hold a certain amount of water. You can think of this as water molecules being dissolved in air. The hotter the air, the more water it can hold. When air is holding as much water as it can at a certain temperature, we say that the air is at 100% relative humidity and saturated with water vapor. If the air becomes colder, some of the water condenses into droplets. We see dew on leaves, grass, and other surfaces in the morning because the temperature has fallen during the night. The air cannot hold as much water as it did during the hotter hours in the day. Much of the dew evaporates once the temperature increases the following day.

A common misconception is that for the air outside to be at 100% relative humidity, it must be raining. It is possible that there are no visible water droplets (rain), and the air simply has the maximum amount of individual water molecules dissolved in it. If you put a cup of water outside on a very humid day without rain, it will not evaporate.

Answer 6:

A great question! There are a number reasons why water may condense after evaporating; it comes down to the right environmental conditions and competition between different energetics, specifically surface energy and volumetric energy (the technical term is Gibbs free energy; it is a kind of potential energy that can be used to do work- you can think of it as extra energy a system has to expend).

Perhaps the example you are most familiar with is water condensation on the side of a container. We'll begin with what we know about systems in bulk (i.e., where the influence of surfaces doesn't matter).

In general, we are familiar with the three (most common) phases of H2O as gas (vapor), liquid (water), and solid (ice). We know that water freezes at 0 degrees C and boils at 100 degrees C. But this turns out to be partially true. You can actually change the temperature at which water boils or freezes if you alter the pressure. The stability of each phase for a particular temperature and pressure is captured in a graph known as a (single-component) phase diagram. (There are phase diagrams for systems with more than one species, but that's a story for another day.) The phase diagram for water is shown here . The familiar temperatures that we recognize as the melting (m.p.) and boiling point (b.p.) of water occur at a pressure of 1 atmosphere, which occurs at sea level where we live. As you can see there is a whole range of temperatures and pressures (blue area) that water can exist. Along the lines between each colored area, you can have a coexistence of the two neighboring phases (e.g., a coexistence of ice and water, vapor and water, or ice and vapor). Where these lines intersect, you can have a coexistence of all three phases.

But this assumes water in its bulk (i.e., no surfaces). This is clearly not the whole picture- there are surfaces everywhere; that is, everything is finite in volume, so surfaces must exist. And surfaces can have a huge influence on material properties and phenomena if your system is small enough (like a water droplet for instance). Associated with every surface is a surface energy. What is surface energy? There are several perspectives on this. The easiest I find to understand is to picture bonding. Most materials in this world involve some kind of bonding between species. This is because it is more energetically favorable to exist as a material with bonding than as an isolated atom or molecule. Species in the bulk of the material have all the bonding they need to be in the most stable configuration. However, species on the surface only have some of the bonding satisfied; there are so-called dangling bonds that exist where the surface is. This is best seen with a picture . The green and blue atoms away from the surface have all the bonding satisfied, but those at the surface do not; the dangling bonds are portrayed with black arrows.

Read the rest of this answer in the text below under "Answer 7"

Answer 7:

There is a constant competition of energetics involving the Gibbs free energy, temperature, pressure, and in this case, surface energy. The goal is to have the lowest free energy (i.e., most stable configuration of species), and the phase that satisfies this is what exists at those particular conditions.

The process of condensation can be considered a subset of processes that fall under nucleation and growth. Nucleation on a surface or defect is known as heterogeneous nucleation (the prefix 'hetero-' meaning different). Condensation specifically applies the phase transformation going from a vapor to liquid phase with any surface. It turns out that have a surface makes it more favorable to condensate water (also known as nucleating water droplets) because condensation lowers the free energy. This is also what happens when clouds form. In the air are lots of particulate matter (e.g., dust, soot, clay) that act as a "surface." Condensation on these particles is energetically more favorable than the particles and water vapor existing separately, as it is (partially) because you are covering up the surface energy of particle. The system of the water and particle is also larger in volume, which offsets the surface energy of the water itself. Of course, this happens within the favorable range of temperatures and pressure that this is allowed.

This same thermodynamics behind condensation (i.e., nucleation of water molecules on a surface) is also the same as when you grow your own rock candy, when you stick your finger in some soda and see the carbon dioxide bubbles gather around it, or how growing any crystalline material used commercially is often started.

In short, it comes down to minimizing this total Gibbs free energy of the system, which involves competition with the surface energies between the vapor/surface, liquid/surface, and liquid/vapor. In this case, the surface is the container the water condensation forms on. Additionally, any defects on the surface make it even easier to nucleate.

This is not really related, but is cool and is something you could try for yourself. There is a phenomenon called supercooling where you lower the temperature past the freezing point without your liquid or gas becoming solid. This can only happen without any defects or impurities (e.g., nucleation centers for ice crystals to form). You can see it in action with this video here . Even after taking it out of the freezer, the water remains a liquid until it is poured out. Upon pouring it out, this metastable state is perturbed and becomes like a slush before melting again. And this can be explained with the snapshot of thermodynamics we just learned!

Hope this helps!

Click Here to return to the search form.

University of California, Santa Barbara Materials Research Laboratory National Science Foundation
This program is co-sponsored by the National Science Foundation and UCSB School-University Partnerships
Copyright © 2020 The Regents of the University of California,
All Rights Reserved.
UCSB Terms of Use