1. Phases of Matter Phase Changes Heating and Cooling Curves
Matter and its Changes
Phases of Matter
Phase Changes
Heating and Cooling Curves

2. Phases of Matter
A phase is a state of existence, a description of how the atoms or molecules of a pure substance are attached to each other.
Chemistry recognizes three (3) common phases: solid, liquid, gas.
Each of the phases has its own characteristics…

3. Solids
Individual particles of the substance are held tightly in place by connections to many other particles
Explains why solids have a rigid shape, definite volume, are typically crystalline, and have all free surfaces.

4. Liquids
Connections between particles are flexible and may be broken. However, as one connection is broken, another connection between the particle and another particle will form.
Explains why liquids will take the shape of their container (flexible connections), still have a definite volume, and only have one free surface.
Also explains why liquids can be poured. (fluid)

5. Gases There are no connections between individual atoms or molecules.
Gaseous systems do not have a definite shape, do not have a definite volume, and have no free surfaces.
Particles are in constant motion and a gas will expand to fill all available space.

6. Phase Changes
This is the description of an event where a pure substance in one state of existence is changed to a different state of existence.
Solid  Liquid is called “Melting”
Liquid  Solid is called “Freezing”
Liquid  Gas is “Evaporation” or “Boiling”
Gas  Liquid is called “Condensation”

7. How Phase Changes Occur
Phase changes will occur when a sufficient quantity of heat has either been added (for melting and boiling) or removed (for freezing and condensation).
Adding heat serves to weaken and/or break the connections between the particles.
Removing heat serves to allow those connections to re-form.

8. Equations of Phase Changes
Solid + heat  Liquid (melting)
Liquid  Solid + heat (freezing)
Liquid + heat  Gas (boiling)
Gas  Liquid + heat (condensation)
Notice that we do not use subtraction signs in chemistry equations describing events.

9. Melting Point A physical property.
Defined as the temperature at which a pure substance will change from solid to liquid.
Is unique for all pure substances.
Is dependent upon the number, type, and strength of the connections between the individual particles of a solid.
The stronger the connections, the higher the melting point.
Is exactly the same temperature as freezing point.

10. Boiling Point Is a physical property.
Is always higher than melting point.
Is unique for all pure substances.
Is defined as the temperature at which a pure liquid will change from liquid to gas.
Also dependent upon the nature of the connections between the particles of the substance.
Is exactly equal to the temperature at which condensation occurs.

11. Heating and Cooling Curves
Graphical representations of how the temperature of a system changes as heat is added or removed through phase changes.
It is observed that the temperature of a system remains constant during any phase change, even though heat is still being added or removed.

12. Heating Curve for Melting
Temp. L
--- melting  S
M.P. L S
Heat added 
Notice how the temperature remains constant during the phase change.

13. Cooling Curve for Condensation
Temp.
--- condensation  L ? G L
--- Heat removed 
Again, notice how the temperature remains constant during the phase change. Also notice that the “x” axis is measuring “heat removed” as the phase change is “downwards”.

14. An Overall Heating Curve
Temp. F 5 4 D E 3 B 2 C 1 A
-- Heat added 

15. A Visual Slide of Phases

16. Sublimation This is a “somewhat exotic” phase change.
In this process, a solid is changed directly to a gas, by-passing the liquid phase.
Only a few substances will do this, examples are dry ice (which is actually solid carbon dioxide), the element iodine, and the element sulfur.

17. Heating curve for Sublimation
temp G
sublimation S G S
Heat added

18. Some final thoughts…
Remember that phase changes are physical changes. You will still have the “same stuff”.
Key thought is that the temperature of the system will remain constant while the phase change is occurring – this is why the temperature stayed the same for so long in your lab.
A final note…there is an opposite process to sublimation. In a very few cases, a gas can be converted directly back to a solid – this change is called deposition.

19. Part II
Liquefaction of Gases

20. What is This ???
This is the process by which a gas is converted into a liquid.

21. How do we do This???
The most easily understood method is to simply remove heat….
The gas particles will slow down as the temperature decreases and the inter-molecular forces will eventually bring the particles together as droplets of a liquid…
But there is another way…..

22. The Use of Pressure
A mechanical system can be constructed that will use pressure to reduce the volume of the system.
This will force the particles of the gas closer together and eventually the intermolecular forces will pull the particles of the gas together as a liquid..
But there is a problem with this system…

23. Pressure vs Temperature
According to Lussac’s Law, as the pressure of a system increases, the temperature of that system will also increase. This increase in temperature will tend to keep the particles moving at high speed, preventing the IMF’s from causing the condensation that you are trying to obtain. P T

24. Commercial Liquefaction of Gases
The process actually uses a system that can be compressed and then expanded a number of times.
First the gas is compressed and additional machinery is used to remove the heat that is produced.
Then the system is allowed to expand. This causes the temperature of the system to decrease (it is called a Joule-Thomson expansion), slowing the particles down even more.
The cycle of compression with refrigeration and expansion is repeated as needed.
This is how we obtain liquid nitrogen, liquid oxygen, etc.

25. Some Vocabulary Terms
Critical Temperature: The highest temperature at which a specific gas can be converted into a liquid using only pressure to cause compression.
The symbol for this measurement is Tcrit
The gas cannot be liquefied at a temperature above this point by pressure alone. (You could still use the commercial process.)

26. Definition
Critical Pressure: This is the pressure that will be needed to convert 1 mole of a gas into liquid at critical temperature.
The symbol for this measurement is Pcrit

27. Next…
Critical Volume: This is the volume occupied by 1 mole of a specific gas at the critical temperature and the critical pressure.
The symbol for this measurement is Vcrit

28. A Definition of Vapor Pressure
Technically, vapor pressure is defined as the pressure of a gas that is contained within a closed system and in equilibrium with its liquid phase.
Picture it this way….

29. Since a vapor is tiny particles, constantly moving, colliding elastically, the particles exert pressure on the surface of the liquid. This pressure is called vapor pressure.
Vapor
Liquid

30. Another discussion of Vapor Pressure
We know that it is the nature of a liquid to evaporate – it increases the entropy – one of nature’s tendencies.
In order for particles of a liquid to evaporate, they must gain enough energy to “jump out of the liquid” and exist as individual gas particles.

31. Continuing
What you must also understand is that the energy gained is needed to enable the liquid particles to “fight their way through” the pressure that is keeping them in the liquid phase.
That pressure is typically the atmosphere.
You should envision it like this…

32. Atmospheric pressure keeping the particles of the substance in the liquid phase.
Gas particles that have “escaped”
Liquid

33. Another Take on Vapor Pressure
Vapor pressure is also defined as the “strength” or energy of the particles in a liquid phase.
This definition leads to a formal definition of boiling and of boiling point.

34. Boiling
The formal definition of boiling is the phase change from liquid to vapor that occurs when the vapor pressure of the liquid is equal to prevailing atmospheric pressure.
In other words, boiling cannot occur until the liquid particles have gained enough energy to fight their way through the atmosphere that is pushing them back into the liquid.

35. Boiling Point
The temperature at which a specific liquid will boil given a prevailing atmospheric pressure exactly equal to 1 atm.
In other words, the temperature at which the particles of a liquid have gained enough energy to make their vapor pressure exactly 760 mm Hg.
Since different substances gain energy differently, we end up with unique boiling points.

36. Phase Diagrams
A phase diagram is a graphical presentation that plots phase of a pure substance as functions of pressure and temperature.
These diagrams have a specific pattern and the questions about phase diagrams are rather predictable.

37. A Typical Phase Diagram

38. So what can you tell with one of these Phase Diagrams??

39. Phases… P T S L G

40. The conditions of the triple point – that combination of temperature and pressure where a pure substance exists in all three phases at the same time. P T S L G

41. You can determine which phase of that pure substance is the most dense.L G
A vertical arrow drawn upward from the triple point will point to the most dense phase.

42. You can use a phase diagram to predict events that will occur if the pressure is changed at a specific temperature.
You can use the phase diagram to predict events that will occur if the temperature is changed at a specific pressure. P T S L G

43. and you can predict the Boiling Point
Remember that boiling point is defined as the temperature at which liquid turns to gas at 1.0 atm.
Therefore, the boiling point is the x-intercept value that intersects a line from the pressure axis to the liquid-gas line at a pressure of 1.0 atm. P T S L G

44. P T S L
1.0 atm G
This would be the boiling point of this substance