1. Chapter 12
States of Matter

2. Kinetic Molecular Theory
Composition and structure (types of atoms & arrangments) determine chem properties of matter
Also affect physical properties
Kinetic Molecular Theory – describes behavior of matter in terms of particles in motion
Kinetic is Greek for “to move”
Ludwig Boltman & James Maxwell proposed model to explain properties of gases

3. Kinetic Molecular Theory
Several assumptions are made by the model:
Particle Size
Gases consist of small particles separated by empty space.
Volume of particles are small compared to volume of empty space
No significant attractive/repulsive forces since particles are so far apart
Particle Motion
Particle Energy

4. Kinetic Molecular Theory
Several assumptions are made by the model:
Particle Size
Particle Motion
Particles are in constant, random motion that move in straight lines until they collide
Collisions between gas particles are elastic. No kinetic energy is lost (may be transferred).
Particle Energy

5. Kinetic Molecular Theory
Several assumptions are made by the model:
Particle Size
Particle Motion
Particle Energy
Mass and velocity determine kinetic energy
𝐾𝐸= 1 2 𝑚 𝑣 2
m=mass
v=velocity (speed and direction)
Temperature measure of average KE in a sample of matter

6. Behavior of Gases KMT explains behavior of gases
Constant motion of particles allows gas to expand and fill container
Low Density (𝑑= 𝑚 𝑣𝑜𝑙 )
Gas with lower density => fewer molecules than another element in the same volume
Compression/Expansion
Compression – reduce volume
Air is compressible
Expansion – larger volume

7. Behavior of Gases Diffusion & Effusion:
Since gas particles are not attracted to one another, they easily pass by each other
When gas flows from one space to another space already occupied by a gas, the gases will mix until evenly distributed (diffusion)
Diffusion – movement of one material through another
Particles diffuse from area of high concentration to low concentration
Ex: food coloring in water, smell when cooking, etc.

8. Behavior of Gases Diffusion & Effusion ~ cont.:
Effusion – gas escapes through tiny opening
Graham’s Law of Effusion –
𝑅𝑎𝑡𝑒 𝑜𝑓 𝑒𝑓𝑓𝑢𝑠𝑖𝑜𝑛 𝛼 1 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠
𝛼 means “inversely proportional to”

9. Behavior of Gases
Heavier particles diffuse slower than lighter particles
Therefore, Graham’s Law can also set up a proportion to compare diffusion rates of 2 gases:
𝑅𝑎𝑡𝑒 𝐴 𝑅𝑎𝑡𝑒 𝐵 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐵 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐴

10. Example 1
Ammonia has molar mass = 17 g/mol, hydrogen chloride has molar mass of 36.5 g/mol. What is ratio of diffusion rates?
𝑅𝑎𝑡𝑒 𝐴 𝑅𝑎𝑡𝑒 𝐵 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐵 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐴
𝑅𝑎𝑡𝑒 𝑁𝐻3 𝑅𝑎𝑡𝑒 𝐻𝐶𝑙 = 𝑔/𝑚𝑜𝑙 17 𝑔/𝑚𝑜𝑙 =1.47
**On the right track as HCl is about twice as massive as ammonia!

11. Example 2
Calculate ratio of effusion rates for nitrogen (N2) and neon (Ne). Soln: 0.849

12. Gas Pressure Pressure – force per unit area
Air Pressure (atmospheric) – pressure is exerted in all directions and varies in different locations due to gravity changes (at different elevations).
P is lower at higher altitudes
At sea level, 𝑃=1 𝑘𝑔 𝑐𝑚 2

13. Gas Pressure Measuring Air Pressure:
Barometer – instrument used to measure atmospheric pressure
At sea level, ~760 mmHg
Changes in air temp or humidity changes pressure!
Manometer – instrument used to measure gas pressure in closed container

14. Units of Pressure Units of Pressure: SI unit is a pascal (Pa)
1 𝑃𝑎=1 𝑁 𝑚 2
**1 𝑁= 𝑘𝑔−𝑚 𝑠 2
At sea level, average air pressure is kPa when 𝑇=0℃
1 𝑎𝑡𝑚=760 𝑚𝑚𝐻𝑔=760 𝑡𝑜𝑟𝑟=101.3 𝑘𝑃𝑎

15. Comparison of Pressure Units
Number Equivalent to 1 atm
Number Equivalent to 1 kPa
Kilopascal (kPa)
101.3 kPa --
Atmosphere (atm)
atm
Millimeters of Mercury (mmHg)
760 mmHg
7.501 mm Hg
Torr
760 Torr
7.501 torr
Pounds per sq. inch (psi)
14.7 psi
0.145 psi
Bar
1.01 bar
100 kPa

16. Dalton’s Law of Partial Pressure
Dalton’s Law of Partial Pressures – Total pressure of mixture of gases = sum of the pressures of all gases in mixture
𝑃 𝑇𝑜𝑡𝑎𝑙 = 𝑃 1 + 𝑃 2 + 𝑃 3 +…+ 𝑃 𝑛
Each gas in a mixture exerts pressure independently of other gasses

17. Example 3
A mixture of O2 gas, CO2 and N2 has a total pressure of 0.97 atm. What is the partial pressure of O2 if the partial pressure of CO2=0.70 atm and N2=0.12 atm? 0.97 𝑎𝑡𝑚=0.7𝑎𝑡𝑚+0.12 𝑎𝑡𝑚+ 𝑃 𝑂 2 𝑃 𝑂 2 =0.15 𝑎𝑡𝑚

18. Chapter 12.2
Forces of Attraction

19. Intermolecular Forces
Intramolecular Forces:
Attractive forces that hold particles together in ionic, metallic, and covalent bonds
“intra-” means forces within
Intermolecular Forces
The forces occur between or among particles
3 types:
Dispersion Forces: weak forces that result in temporary shifts in density of electron clouds
Dipole-Dipole – attraction between oppositely charged regions of polar molecules
Hyrdrogen Bonding – (type of dipole-dipole) bond that exists between H and one of the following: F, O, N

20. Dispersion Forces Weak forces
Sometimes called London forces after physicist Fritz London
Electrons in electron cloud are in constant motion
When 2 electron clouds come in contact, they repel
For q quick moment, the electron density is greater in one region forming a temporary dipole
All particles consist of dispersion forces
As size of particle increase, dispersion forces increase too (stronger)

21. Dipole-Dipole Forces
Polar molecules contain permanent dipoles (some regions are always partially pos/neg)
Neighboring molecules orient themselves so oppositely charged particles will align

22. Hydrogen Bonds Typically stronger than dipole-dipole and dispersion
H bonding only occurs with F, O, N (FON or “phone”)

23. Comparison

24. Chapter 12.3
Liquids and Solids

25. Liquids Recall, (Chap 3) Take shape of container!
Forces of attraction keep molecules closely packed in fixed volume, but not in fixed position.
Density and Compression:
Liquids are much denser than gases - stronger intermolecular forces holds particles together.
Large amounts of pressure must be applied to compress liquids to very small amounts.
Liquids considered incompressible

26. Liquids
Fluidity - ability to flow and diffuse; liquids and gases are fluids.
Liquids diffuse slower than gases at the same temp due to intermolecular forces interfering with flow
Viscosity - measure of resistance of a liquid to flow
Determined by type of intermolecular forces, size and shape of particles, and temp
Particles in liquid are close enough for attractive forces to slow their movement as they flow past one another
The stronger the intermolecular attractive forces, the higher the viscosity

27. Liquids Viscosity~Cont. Larger molecules create greater viscosity.
Long chains of molecules result in higher viscosity.
Increasing the temp decreases viscosity - adds energy that allows molecules to overcome intermolecular forces and flow more freely.

28. Liquids
Surface Tension - energy required to increase the surface area of a liquid by a given amount.
Particles in middle can be attracted to particles above, below and on either side
No attractive forces above for particles on surface to balance forces from below
Net forces is pulled down => surface pulled tight
Stronger attractive forces means stronger surface tension
**Soaps/detergents decrease surface tension by disrupting H-Bonds. When bond is broken, water spreads out allowing dirt to be carried away!
Surfactants- compounds that lower the surface tension of water.

29. Liquids
**When water placed in glass tube, the surface of the water is not straight (meniscus)
Cohesion - force of attraction between identical molecules
Adhesion - force of attraction between molecules that are different
Capillary action - upward movement of liquid into a narrow cylinder, or capillary tube (how paper towels work!)

30. Solids Recall: Definite volume Definite shape
Most solids are more dense than liquids.
Ice is not more dense than water.
Particles in solid ARE in constant motion!
Particles in a solid vibrate in a fixed position
To be solid vs. liquid at room temp, there must be strong attractive intermolecular forces
These forces limit motion of particles in vibration creating order

31. Solids
Crystalline Solids - solids with atoms, ions, or molecules arranged in orderly, geometric shape
Location of particles is are represented by points on crystal lattice frame work

32. Solids
Unit Cell – smallest arrangement of atoms in crystal lattice that has same symmetry as whole crystal (building block)
Crystal shapes differ since surfaces/faces of unit cells do not always meet at right angles

33. Solids 5 Categories of Crystalline Solids: Molecular Solids
Solids held together by dispersion, dipole-dipole or H-Bonding
Most not solid at room temp
Poor conductors of heat and electricity
Covalent Network Solids
Form multiple covalent bonds (C, Si, etc)
Ex: carbon forms diamonds
Ionic Solids
Metallic Solids
Amorphous Solids

34. Solids 5 Categories of Crystalline Solids: Ionic Solids
Each ion is surrounded by ions of opposite charge
High MP and Hardness
Strong but brittle
Metallic Solids
Positive metal surrounded by sea of electrons
Malleable and ductile
Good conductors of heat and electricity
Amorphous Solids
Particles not arranged in regular, repeating pattern (no crystals)
Generally form when molten material cools too quickly to allow crystals to form
Ex: glass, rubber, plastics

35. Chapter 12.4
Phase Changes

36. Phase Changes Phase changes require energy!
When Energy is added/removed from system, one phase can change into another

37. Phase Changes Melting: What happens when you put ice cubes in water?
Water high temp & heat flows from water to ice!
Heat is transfer of energy from object with higher temp to object with lower temp
At ice’s MP, the E ice absorbs disrupts the H-bonds that hold water molecule together in ice crystal
When enough E is absorbed, the H-bonds break and move apart to enter liquid phase
Melting Point – temp in which forces holding crystal lattice together are broken & become liquid

38. Phase Changes Vaporization:
Process by which liquid changes from liquid to a gas/vapor
When ice melts, temp of ice and water remain constant until ice is melted completely
If heat still being added, THEN temp will increase
**Vapor – gas phase of substance that is usually a liquid at room temp
If input of E is gradual, molecules tend to escape from surface of liquid. When vaporization occurs only at surface of a liquid, you get evaporation.

39. Phase Changes

40. Phase Changes
Vapor Pressure – pressure exerted by a vapor over a liquid
Boiling Point – temp where VP of liquid = the atmospheric pressure
At BP, bubbles begin to form indicating an increase in E
Sublimation – process of a solid changing directly to a gas
Ex: dry ice, moth balls, air fresheners, etc.

41. Phase Changes Phase changes that Release E:
Freezing Point – temp where liquid is converted into crystalline solid
Heat is removed from system, KE decreases & velocity decreases
H-bonding increases creating a solid
Condensation – process where gas/vapor becomes a liquid
Vapor loses E, velocity decreases, H-bonding increase
Ex: condensation on glass, dew on grass, etc.

42. Phase Changes
Deposition – process of gas/vapor changing to solids without becoming a liquid
Ex: frost

43. Phase Changes
Phase Change Diagrams – graphs of P vs. T that show phase of different substances under different conditions
temp & pressure control phase of substance
Pts on curve is where 2 phases coexist
Triple Point – point that represents temp and pressure where all 3 phases exist
Critical Point – point of critical temp and pressure water can NOT exist as liquid

44. Phase Changes