There are two ways for a liquid to vaporize (to go from a liquid state to a gas state): evaporation and boiling. The primary difference between evaporation and boiling is where vaporization occurs. The temperature of a substance is a measure of the average kinetic energy of the molecules in the substance. But even if the average kinetic energy of the molecules in a substance is quite low, there will always be some high-energy molecules. In evaporation, those high-energy molecules have a chance of breaking away from the liquid state and entering a gas state only when they are near the surface of the liquid. What happens to the high-energy molecules in the center of the liquid?
The blue molecule has twenty-five times as much kinetic energy as the other molecules. But when the blue molecule is at the bottom of the liquid, it has nowhere to go. For all its energy and speed, the blue molecule simply bounces back and forth between collisions with other molecules. And in each collision, the blue molecule loses energy to the slower moving molecules around it. The only way the blue molecule can enter a gas state is if it is at the surface of the liquid… or if there are enough high-energy molecules to form a bubble.
To form a bubble beneath the surface of a liquid, there needs to be enough blue molecules moving fast enough to lift the liquid sitting on top of them. This means that the vapor pressure created by the blue molecules inside of the bubble must be at least as great as the air pressure pushing down on the liquid. (Remember, pressure is created in a fluid by the collision of molecules. The faster those molecules are moving, the more force in their collisions and the greater the pressure.)
According to the ideal gas law: PV = nRT, where P is the pressure, V is the volume, n is Avogadro’s number, R is the gas constant (8.314472 J·K-1·mol-1), and T is the temperature in kelvins (K). Now take a look at that formula again and think about what it means. If the temperature of a gas increases, then the pressure of the gas and/or the volume of the gas must increase as well. A classic way to visualize the ideal gas law is in a piston. Imagine water molecules in a gas state trapped beneath the weight of liquid water and the air pressure pushing down on the liquid water. As you increase the temperature, see what happens to the volume of the water vapor “bubble.”
At higher temperatures, the pressure from the collision of water molecules in the gas state increases until the vapor pressure is great enough to lift the weight of the liquid water and a bubble begins to form. After that, any increase in temperature simply increases the size of the bubble. This is how the pistons inside of an internal combustion engine work. The energy from the up-and-down motion of the piston is transferred to an axle that turns the wheels of a car.
At 100 °C, the average speed of a water molecule is approximately 660 m/s (≈1500 mph). At this speed, the vapor pressure generated by the collision of water molecules will equal standard atmospheric pressure, and bubbles of water vapor will begin to form in liquid water. Below this speed, bubbles of water vapor will not have the vapor pressure to “inflate” underwater. Once those bubbles rise to the surface, the water molecules inside of the bubbles will then break away from the liquid and form water vapor (or steam) in the air. Now, instead of vaporization only occurring at the surface of the liquid, vaporization is also occurring within the liquid. This process is called boiling and the temperature at which it occurs is called the boiling point.
The boiling point depends on the properties of the molecules in the liquid and the air pressure pushing down on the liquid. For example, the city of Quito in Ecuador is 2800 m above sea level and it has an air pressure that is only ≈70% of standard atmospheric pressure. That means that, in Quito, water molecules will not need to be moving as fast to generate the vapor pressure needed to form bubbles of water vapor in liquid water, and water will boil at approximately 90 °C. Meanwhile, pressure cookers are designed to increase the pressure on a liquid, and most pressure cookers can increase the pressure up to twice standard atmospheric pressure. At those pressures, water molecules will have to move much faster to generate the vapor pressure needed to form bubbles of water vapor in liquid water, and water will only boil at temperatures over 120 °C.
When you add heat (energy) to water, the temperature of the water goes up. Makes sense, right? Except, that does not always happen. If the heat you are adding to the water goes into increasing the kinetic energy of the water molecules (increasing their speed), then the temperature will go up. But molecules have other forms of energy that the heat can go into. One of them is electric potential energy.
Something curious happens when liquid water reaches its boiling point. The temperature of the water stops going up even though you are still pumping heat into it. So, where does all of this heat go? At 100 °C (the boiling point of water at standard atmospheric pressure), water molecules in the liquid state have a density of 0.9584 g/cm3. This means that there are 3.21 × 1022 water molecules packed into 1 cm3 and the average distance between water molecules is 3.91 × 10-8 cm. When that liquid water vaporizes and the water molecules enter the gas state, the volume of the water undergoes a massive expansion: the density drops to 0.0006 g/cm3 and the average distance between water molecules grows to 4.6 × 10-7 cm. In the transition from the liquid state to the gas state at the boiling point, the average distance between water molecules becomes almost twelve times greater.
As you have learned, pulling molecules apart that are held together by intermolecular forces takes energy. In the case of water molecules, which are strongly attracted to each other, it takes a lot of energy. That energy can come from the kinetic energy of the water molecules themselves, but this would mean that the water would cool as it expands from a liquid to a gas. And if the water cools in order to expand, then its temperature would drop below its boiling point.
If you were to heat liquid water to 100 °C and then turn off the heat, the liquid water will not vaporize and water molecules will not enter the gas state (except through evaporation at the surface of the liquid). Heating liquid water to 100 °C is not enough. To convert liquid water into water vapor (steam), you must continue to supply heat to the water so that its molecules have the energy to overcome the intermolecular forces holding them together, and the water can expand into a gas. It takes 420 J of energy to heat 1 g of liquid water from 0 °C to 100 °C. It takes 2260 J of energy to pull apart the water molecules in 1 g of liquid water and to transition them into the gas state.
The energy required to separate water molecules and to transition them from a liquid state into a gas state is called the heat of vaporization. The heat of vaporization for water at 100 °C and standard atmospheric pressure is 2.26 kJ/g. The reason why the temperature of liquid water plateaus (stops going up) once it reaches its boiling point (even though heat is still being pumped into it) is because, once water begins to boil, any additional heat added to the liquid will go into converting liquid water into water vapor (steam), and not into increasing the speed of the water molecules themselves. The temperature of the water will not start climbing again, going above the boiling point of water, until all of the liquid water has been boiled off and converted into a gas.
The mole (mol) is a unit of measurement used to express amounts of a substance. It is especially useful in chemistry when you want to know the number of molecules in a substance. One mole of water contains 6.022 × 1023 water molecules. One mole of ethanol contains 6.022 × 1023 ethanol molecules. One molecule of water has a mass of 18 u; one mole of water has a mass of 18 g. One molecule of ethanol has a mass of 46 u; one mole of ethanol has a mass of 46 g. (If the number 6.022 × 1023 seems familiar, it is because it is Avogadro’s number.) So while we can express the heat of vaporization in terms of grams (2.26 kJ/g), sometimes it is convenient to also express the heat of vaporization in terms of moles (40.65 kJ/mol).
The heat of vaporization is stored in the water molecules in the gas state as electric potential energy. When steam condenses back into a liquid state, that electric potential energy is turned back into kinetic energy (as the water molecules are pulled closer together, the force of attraction between molecules causes the molecules to accelerate towards each other, increasing their speeds). This is why the burns caused by 100 °C steam are often much more severe than the burns caused by 100 °C liquid water. Much of the energy that is released back as heat when water condenses is transferred to your skin.