1. Phases of Matter

2. Phases
Solid
Liquid
Gas
Plasma

3. Phases of Matter IMF’s explain the phases of matter Solid
Particles move relatively slow (basically vibrating in fixed position)
IMF’s hold them together
Liquid
Particles move more (able to flow past each other)
IMF’s still hold together but not as tightly
Gas
Particle move quickly (each in rapid, random motion)
Little to no IMF’s (we usually assume none)

4. Intermolecular forces are very important.
Intermolecular forces are of little significance
Intermolecular forces must be considered.

5. Solid Particles relatively close together Retains volume and shape
Relatively high densities
Hard to compress
Does not flow easily
Types
Crystalline- regular order/ pattern to particles
Amorphous- no regular pattern to particles

6. Liquid Particles more spaced apart Retains volume but not shape
Intermediate densities
Hard to compress
Flows easily (Fluidity)
Slow flow faster than others
Viscosity- measure of the resistance to flow in a liquid
Stronger IMF’s lead to higher viscosity
Can diffuse- liquid molecules spread out through another liquid
Can display surface tension- attraction of molecules at surface to each other (Stronger IMF’s lead to more surface tension)
Can display capillary action- attraction of liquid to solid surface causing it to flow

8. Gas Particles spaced very far apart Does not retain volume or shape
Very low densities
Compressible
Flows easily (fluidity)
Can diffuse and go through effusion
Elastic collisions (no loss of energy during collision)
Ideal vs Real Gases
Ideal- no attraction between particles
Real- gases where particles are attracted to each other

9. Plasma Occurs at very high temperature and pressures Also a fluid
Mix of neutral atoms, free electrons, and ions

10. Density g/cm3 or g/ml (1cm3 = 1 ml) g/L Physical, intensive property
Mass per unit volume of substance (D= m/V)
Units
Solids and Liquids
g/cm3 or g/ml (1cm3 = 1 ml)
Gases
g/L
Density controls placement of fluids and solids
Less dense objects or fluids move to the top
More dense objects or fluids move to the bottom

11. Density (cont) D = m/V Mass Volume Measured on a balance Solid
Regular Shape- can be calculated from other measurements
Irregular Shape- can be found by water displacement method
Liquid and Gas
Can be measured with instruments such as graduated cylinder

12. Heat Amount of energy transferred from one substance to another
Represented by q with units in Joules (J)
When heat transfers, it affects the temperatures of the substances

13. Temperature Measure of the average kinetic energy in a sample
High temperatures mean the particles are moving quickly
Theoretically if the particles weren’t moving at all, the temperature would be 0 Kelvin (absolute zero)
Remember K = °C

14. Heat Flow
Heat will “flow” from the substance with a higher temperature to the substance with a lower temperature
As the faster particles collide with slower particles, the faster ones will slow down and the slower ones will speed up

15. Heat and Temperature Change
When heat transfers, it affects the temperatures of the substances involved in the transfer
How much will the temperature change?
Dependent on
Amount of heat transferred
Mass of the sample
Composition of the sample

16. Amount of Heat Transferred
The more heat transferred, the greater the temperature change
If heat is absorbed by the sample
q is positive
Final temperature will be higher than the initial temperature
If heat is lost by the sample
q is negative
Final temperature will be lower than the initial temperature

17. Mass of Sample
A heat transfer will cause a bigger temperature change to a smaller mass than it will to a larger mass.

18. Composition of the Sample
Different substances absorb/release heat in different ways.
Specific heat (c) – is the amount of heat needed to change 1 gram of a particular substance by 1 °C.
Each type of substance has a different value

19. Specific Heat Equation
q= mcT
q is heat
m is mass
c is specific heat
T is change in temperature
Tf-Ti (final temperature – initial temperature)
Units need to match
For instance, if the specific heat value is given in J/g°C, heat should be in J, mass in grams, and temperature in °C

20. Calorimeters Instrument used to measure heat transfer
Process allows water and another substance to undergo a heat transfer until both are at the same temperature (thermal equilibrium)
This means the heat gained (or lost) by the water equals the heat lost (or gained) by the substance

21. Absorbing Heat
Solid absorbs heat and temperature increases (molecules moving faster)
Reaches a point that movement weakens IMF’s enough to allow flow (melting point)
Heat is still absorbed but temperature does not increase
Liquid absorbs heat and temperature increases (molecules moving faster)
Reaches a point that movement weakens IMF’s enough that they essentially no longer exist (boiling point)
Gas absorbs heat and temperature increases (molecules moving faster)

22. Heating Curve Plateaus at melting point Heat still added
Used to weaken IMF’s
Called Heat of fusion
Heating Curve
Another longer plateau will occur between liquid and gas
Called heat of vaporization

23. Releasing Heat
Heat is released from the gas and temperature decreases (molecules moving slower)
Reaches a point that molecules are close enough for IMF’s to be reestablished (condensation point)
Heat is still released but temperature does not decrease
Heat is released from the liquid and temperature decreases (molecules moving slower)
Reaches a point that molecules are close enough for IMF’s to strengthen (freezing point)
Heat is removed from the solid

24. Heating/Cooling Curve

25. Phase Changes Solid  Liquid Liquid  Gas Solid  Gas
Solid  Liquid = Melting
Liquid  Solid = Freezing
Occurs at melting/ freezing point
Liquid  Gas
Liquid  Gas = Evaporation/Vaporization/Boiling
Gas  Liquid = Condensation
Occurs at boiling/condensation point
Solid  Gas
Solid  Gas = Sublimation
Gas  Solid = Deposition

26. Evaporation or Vaporization
Occurs at temperatures below boiling point
Some molecules have enough energy to escape surface of liquid
Vaporization
Occurs at boiling point
Change to gaseous phase occurs throughout liquid

27. Phase Change (cont)
Freezing point and Melting point
Same thing (occur at same temperature)
Named depending on the direction compound is going
Condensation point and Boiling point
Each substance has its own points and heats (fusion and vaporization)

28. Influencing Points Pressure (mostly sways boiling point)
Same substance
Pressure (mostly sways boiling point)
“Normal” points are points at standard pressure (1atm)
Lower pressures allow particles to spread out more (IMF’s can be overcome at lower temps)
Higher pressures compress molecules (Higher temp needed to overcome IMF’s)
Between different compounds
Strength of forces holding particles together
Metallic Bonds
Ionic Bonds
Covalent Bonds
IMF’s

30. Boiling Point
Vapor pressure- Partial pressure of gas particles of substance over the liquid of that substance
Vapor pressure increases with temperature
More particles have energy to escape surface
Point at which vapor pressure of substance is equal to atmospheric pressure

31. Phase Diagram
Chart for each substance showing the temperature for phase changes according to pressure
Crossing a line indicates a phase change

33. Phase Diagrams (cont) At any pressure, a horizontal line can be drawn.
Temperatures of phase changes are found where lines are crossed
The “normal” points are found by drawing a horizontal line at 1atm of pressure

35. Phase Diagrams (cont) Triple Point Critical Point
Pressure and Temperature where all three phases can be found
Critical Point
Critical Temperature- highest temperature that the liquid phase of a substance can be found
Critical Pressure- pressure at critical point
Beyond this point the liquid and gas phase in indistinguishable (super critical fluid)