Within the liquid, the molecules are still moving with respect to one another, that is why liquids flow. What does that mean, at the molecular level? It means that the molecules are (on average) moving fast enough to break some, but not all of the interactions linking them to their neighbors. But let us consider what happens as we remove more and more energy from the system (by interactions of the molecules with the walls). Now the frequency with which molecules have sufficient energy to break the interactions between them decreases; interactions become more stable. Once most interactions are stable, the substance becomes a solid. The temperature at which the material goes from solid to liquid is termed the melting (or freezing) point. For water at atmospheric pressure, this is 0 ºC (or 273.15 K). Just like the boiling/condensation point, the temperature does not change appreciably until all the liquid has solidified into ice.

5.1 Systems
5.2 Temperature
5.3 Vibrations
5.4 Phase changes
5.5 Thermodynamics
5.6 Phases, again

What happens when water (or any other liquid) is cooled and eventually freezes is determined by molecular shape and the geometry of the interactions between molecules. In the case of frozen water (ice), there are in fact more than 15 types of arrangements of the molecules, ranging from amorphous to various types of “crystalline” ice.

In amorphous ice the molecules occupy positions that are more or less random with respect to their neighbors; in contrast, the molecules in crystalline ice have very specific orientations to one another. The most common form of ice, and the one we are most familiar with is known as Ice Ih: the water molecules are organized in a hexagonal, three-dimensional array. Each molecule is linked to four neighboring molecules through H-bonds. This molecular level structure is reflected at the macroscopic level - which is why snowflakes are hexagonal. Because of the bonds between them, the molecules ?can no longer move with respect to one another - ice is solid and retains it shape, at both the visible and the invisible (molecular) level. That said, since we are not at absolute zero (0 K, -273.15 ºC), the molecules are still vibrating in place.

Now let us go backward, and transfer energy from the surroundings into the system, for example by heating our container. The energy is stored in the water (system) by increasing molecular vibrations. Eventually the molecules will vibrate so vigorously that the hydrogen bonds that are holding the molecules in place will be broken and the molecules will become free to move relative to each other. The ice will melt. At this temperature (0 ºC, 273.15 K), all the energy entering the system will be used to overcome intermolecular attractions (that is to break bonds, rather than increase the speed of molecular motion) – if the system is well mixed, the temperature will stay at 0ºC until all of the ice has melted, after which the temperature will start to rise again as the water molecules, now free to move relative to each other, increase in kinetic energy.

Because of the arrangement of water molecules in hexagonal ice (Ice Ih), the hexagonal “cages” of water molecules within the crystal have empty space within them. As the hydrogen bonds break, some of the water molecules can now approach closer to each other - filling in the open cages. The structure of the ice collapses in on itself. This open network of molecules, which is not present in liquid water, means that Ih ice is less dense than liquid water – Ih ice floats on liquid water. We don’t think much of this commonplace observation, but it is quite rare for a solid to be less dense than the corresponding liquid. More typically, most materials expand when heated (particularly gases - but also liquids and solids) as a consequence of the increased kinetic energy making the particles vibrate more vigorously - which means that they take up more space.

Open versus closed systems

In our discussion, our container of gas (water vapor) was our system, that is the part of the universe we are observing. It is separated from the rest of the universe (its surroundings) by the walls of the container (the boundary). When we remove energy from the system or add energy to it, that energy goes to or comes from the surroundings. Our system is certainly not an isolated system; in an isolated system neither energy nor matter move between the system and the surroundings. In practice it is difficult to construct a perfectly isolated system - although an insulated (styrofoam) coffee cup with a lid on it is not a bad approximation. We can also distinguish between open and closed systems; in an open system, both matter and energy can enter or leave (but we keep track of both), while in a closed system, the amount of matter is constant, but energy can enter or leave. Whenever we look at a system, our first task is to decide whether the system is isolated, open or closed. All biological systems are open - that is both energy and matter are being exchanged with the surroundings – in the absence of such an exchange, a biological system will (eventually) die [The only exception would be cryptobiotic systems, like the tardigrads mentioned earlier.].

So let us consider a beaker of water without a lid as an open system. As the temperature rises, some of the water molecules will have enough energy to escape from the body of the water. The liquid water will evaporate (that is, change to a gas). Any gases that might be dissolved in the liquid water, such as oxygen (O₂) or nitrogen (N₂) will also move from the liquid to the gaseous phase.

At the boiling point, all the energy being supplied to the system is being used to overcome the intermolecular forces -similar to the way it did at the melting point - but this time the molecules separate from one another completely (although they still collide periodically.) That is: energy is used to over-come attractive forces and the individual molecules fly off into the gas phase where the distances between them they become so great that the attractive forces are insignificant . As the liquid boils the temperature (of the liquid) does not rise until all the liquid has boiled. As the gas molecules fly off, they carry with them some of the system’s energy.

5.1 Systems
5.2 Temperature
5.3 Vibrations
5.4 Phase changes
5.5 Thermodynamics
5.6 Phases, again

Question to answer:

• Begin with an ice cube in a beaker and end with water vapor. Draw a graph of the energy input versus the temperature of the system; is your graph a straight line?
• What would happen to the mass of the beaker + water during this process?
• Can you reproduce the hexagonal symmetry of ice by using a model kit? What property of hydrogen bonds is it that makes the structure so open?
• As the temperature rises in liquid water what do you think happens to the density? Draw a plot of density v temperature for a mass of water beginning at -10 ºC to 50 ºC.
• What happens when the temperature has risen such that the molecules have enough energy to overcome all the intermolecular attractions between them? (Not the covalent bonds - but the attractions between separate molecules)
• During evaporation and boiling do water molecules ever return to the liquid?
• Estimate the temperature at which the bonds within a water molecule break; how does that temperature compare to the boiling point of water?
• Why aren’t they the same temperature?
• How would an open and a closed system differ, for example if you heated them from 30 to 110ºC?
  

Questions to ponder and questions for later:

• Are boiling and evaporation fundamentally different processes.
• Under what conditions does evaporation not occur? What is happening at the molecular level?
• What is in the spaces in the middle of the hexagonal holes in ice?
  What would be the consequences for an closed or an isolated biological system?
• As you heat up a solution of water, predict whether water molecules or dissolved gas molecules will preferentially move from the liquid to the gaseous phase (or will they all move at the same rate?)
• What factors do you think are responsible for “holding” the gas molecules in the water.
• What do you think happens to the density of the gas (in a closed system?) as you increase the temperature?
• What would happen if you captured the gas in a container?
• What would happen if you took that gas in the container and compressed it? (made the volume of the container much smaller)