Boiling water with ice!


Atomic Theory

“All things are made up of atoms.”

States of matter

Solid, liquid, gas, and plasma

State of Solids

  • Has definite shape and volume
  • High density and not very compressible
  • Does not depend on the shape of the container (doesn’t fill it in)

State of Liquids

  • Has a fixed volume
  • Takes the shape of container
  • Less dense than solids
  • Almost incompressible

State of Gases

  • Takes the shape of the container
  • Takes the volume of container (this means that the gas molecules can be spread out evenly in the container or compressed in a smaller container. The volume is when the molecules are evenly distributed.)
  • Can be compressed
  • Gases are in the gaseous state at room temperature
  • Gases have more energy than liquids, and solids.
  • Gases are less dense.

State of Plasma

Plasma is simply an ionized gas, where the gas is charged with free electrons and positive ions because of the amount of energy plasma contains.

Unit 1.5 Vapour Pressure

As you have read from the previous section of phase changes, heat or energy is required for a substance to change states. But where do “boiling points” or “freezing points” come from? What exactly is happening at the boiling point?

Vapour pressure of a liquid is the pressure exerted by the vapour on the liquid at equilibrium. In other words, when a liquid is evaporating and condensing at equilibrium, the vapour pressure is the amount of force exerted by the vapour.

How I like to think of vapour pressure is that if we imagine the molecules at the surface of the liquid having a “push and shove” contest with the particles in the atmosphere.  The greater the vapour pressure, the greater the force the liquid has to “push” into the vapour stage. If the force of the vapour pressure is the same as the atmosphere pressure, this is when dynamic equilibrium occurs with the liquid entering the gas and liquid phase simultaneously.


How to read and interpret a vapour pressure graph: 

The vapour pressure graph above depicts the vapour pressure of various substances over a range of 120 degrees. Each line represents the vapour pressure of each substance at that particular temperature. If we use water as an example, at 80C, water has a vapour pressure of about 50kPa. Recall that phase change occurs when the vapour pressure equals to the atmospheric pressure, thus if the atmospheric pressure were to be lowered to 50kPa, then the water will change states into the gas state.

Since our atmospheric pressure hovers at 101.3kPa, in order for water to enter the gas phase is if the temperature of the water hits 101.3kPa, and this is when water is at 100C.

If we look at ethanol as an example, in order for ethanol to boil is if the temperature hits about 80C.  This is when the ethanol contains enough kinetic energy for the molecules to overcome the atmospheric pressure of 101.3kPa.

At this point, you might be wondering? Well don’t water evaporate at room temperature? Yes, that is correct.  Evaporation and boiling are two similar but different processes. Both evaporation and boiling are under the same category of vaporization, which is changing from a liquid to a gas state.  In evaporation, it tends to occur at the surface of the liquid and this can occur at any temperature, even in cold temperatures while boiling is heating the water to the above the pressure being applied from the outside, such as the atmospheric pressure.

Changing the Boiling point and freezing points 

Boiling occurs when the liquid contains enough kinetic energy to overcome the pressure exerted from the outside, such as the atmospheric pressure. If there is a high atmospheric pressure exerted on a liquid, the energy required to boil will be higher since it will need more energy to enter the gas phase.  However, if the atmospheric pressure is low such as, at the peak of a mountain, the liquid will not need as much energy to boil, thus lowering the boiling point.

Boiling point depression and elevation will depend on the external pressure. In most cases, the pressure of concern will be out atmospheric pressure.  If you think relate this to the vapour pressure example that the pressure of the liquid is having a “push and shove” contest with the atmospheric pressure. If there is a higher amount of force or a lot of particles “sitting” on the surface of the liquid, you will need to “push and shove” harder to over them those extra particles. By having to “push and shove” harder, then you will need more energy, thus elevating the boiling point!

In contrast, if we were to have a LOW atmospheric pressure, there will be less atmospheric pressure, or particles to “push and shove” against the liquid, thus the liquid does not need to push as hard to enter the gas phase, thus lower boiling point. One cool demonstration is boiling water with ice or in a vacuum!  Even though the water is boiling, it does not mean that it is boiling at 100C!