1. X Unit 13
States of Matter

2. Reviewing States of Matter

3. Reviewing States of Matter1) 2) 3)

4. Solid States of Matter Has definite shape & volume
Particles are tightly packed
Can expand when heated

5. Liquid States of Matter
Has constant volume but takes the shape of its container
Fluid
Less closely packed particles than solid particles
Can expand when heated

6. States of Matter
Gas
Expands to fill its container & takes the shape of its container
Fluid
Much less closely packed than solid particles
Expands when heated

7. Forces of Attraction
There are two kinds of attractive forces at the molecular level –
The forces inside a molecule holding the individual atoms together
The forces between molecules holding different molecules together in a sample

8. Intramolecular Forces
“Intra-” prefix = within
The forces inside a molecule holding the individual atoms together
Ex.) Covalent bonds in H2O

9. Demo
What happened to the paper clip when placed in the beaker of water vs. the beaker of acetone? Explain your observations.

10. Intermolecular Forces
“Inter-” prefix = between
Short range forces between molecules in a sample
There are 3 main types of intermolecular forces
Hydrogen bonding
Dipole-dipole forces
London Dispersion forces

11. Dipole-Dipole Forces

12. Dipole-Dipole Forces
Dipole – a molecule or part of a molecule that contains both positively and negatively charged regions
δ+ (partial positive) or δ- (partial negative)
Dipole-Dipole Forces – forces of attraction between POLAR molecules
Dipoles must get close together in correct orientation (positive end must be near negative end)

13. Dipole-Dipole Forces H-F H-F H-F H-F
Dipole-dipole forces will raise melting and boiling points.
A dipole can temporarily attract electrons from another molecule causing an induced dipole.

14. London Dispersion Forces

15. London Dispersion Forces
Intermolecular attractions resulting from the uneven distribution of electrons and the creation of temporary dipoles
Present in all substances (polar molecules, non-polar molecules, and noble gases)
The weakest intermolecular force
Does the number of electrons play a role regarding strength of the attraction??

16. London Dispersion Forces
Electrons are constantly moving around the nucleus therefore electron density can fluctuate
This effect becomes stronger with increasing number of electrons
Example: F2
Br2 I2

17. Hydrogen Bonding

18. Hydrogen Bonding
Hydrogen Bonding – attractive forces in which a hydrogen covalently bonded to a very electronegative atom (F, O, N) is also weakly bonded to an unshared electron pair on another electronegative atom (another F, O, or N atom)
Hydrogen bonds are strong intermolecular forces

19. Hydrogen Bonding
Hydrogen bonding raises melting and boiling points because more energy is required to break the forces between molecules.

20. Hydrogen Bonding in DNA
Hydrogen bonding plays a key role in maintaining the double helix structure of DNA

21. Think Back to the Demo Question… Why did the paperclip float on water but not acetone?

22. Before You Go:
Identify the types of intermolecular forces present in compounds of:
Hydrogen Fluoride
Pentane (C5H12)
Hydrochloric Acid
Ethanol (Ethyl Alcohol)

23. Relative Strength of Intermolecular Forces
(Strongest)
(Weakest)

24. Liquids & Solids
Even though they are more restricted than gas molecules, the molecules of solids and liquids are constantly in motion as well!
(This idea comes from the Kinetic Molecular Theory – we’ll come back to this idea)

25. Liquids Viscosity – a measure of the resistance of a liquid to flow
The particles in a liquid are close enough together that their attractive forces slow their movement as they flow past one another
The stronger the attractive forces (intermolecular forces), the more viscous the liquid is.
As temperature increases, viscosity decreases.

26. In water, this is due mainly to hydrogen bonding!
Liquids
Surface tension – an inward force that tends to minimize the surface area of a liquid
A measure of the inward pull by particles in the interior
The stronger the intermolecular forces, the higher the surface tension
In water, this is due mainly to hydrogen bonding!

27. Liquids
Surfactant – any substance that interferes with the hydrogen bonding between water molecules & reduces surface tension

28. Surfactants used to clean up oil spills as well
Exxon Valdez oil spill in 1989 spilled over 700,000 barrels of oil into the water near Alaska

29. Solids
Crystalline solid – a solid in which the atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure
Unit cell – the smallest arrangement of connected points that can be repeated to form the lattice
A.k.a. The representative part of the whole crystal

30. Crystalline Solids
Pyrite (cubic)
Rutile (tetragonal)

31. Crystalline Solids
Copper sulfate
(triclinic)
Borax (monoclinic)

32. Network Covalent Solids
Atoms that can form multiple covalent bonds
(such as C, Si, and other Group 14 elements)
are able to form network covalent solids.
All atoms in the entire structure are bonded together with covalent chemical bonds.

33. Metallic Solids
Metallic solids – positive metal ions surrounded by a sea of mobile electrons
Mobile electrons make metals malleable and ductile because electrons can shift while still keeping the metal ions bonded in their new places
Metallic solids are good conductors of heat and electricity
Metallic Bonds

34. Amorphous Solids
Amorphous solid – a solid in which the particles are not arranged in a regular, repeating pattern
“Amorphous” = “without shape”
Often form when a molten material cools too quickly to allow enough time for crystals to form
Common examples: glass, rubber, many plastics

35. Kinetic Molecular Theory
Describes the behavior of gases (and solids/liquid to some extent) in terms of particles in motion
Assumptions:
Gas particles have negligible volume compared to the volume of their container
Particles move in constant, random, straight line motion
Particles collide with themselves and walls without losing energy (elastic collisions)
There are no intermolecular forces between gas molecules

36. Kinetic Molecular Theory
If gas particles are always in this constant, random motion, what keeps them going?
ENERGY!!
(Kinetic Energy to be exact)
Temperature is a measure of the average kinetic energy in a substance.
Higher temp. = more kinetic energy = particles move faster!

37. Temperature Celsius (°C) Kelvin (K) 3 Main Temperature Scales
Fahrenheit
Celsius (°C)
°C = K - 273
Kelvin (K)
K = °C + 273

38. Temperature Conversions
Convert the following temperatures:
28 °C = ________ K
200 K = ________ °C
-15 °C = _________ K

39. Gases
Remember:
Gases expand to fill their container & take the shape of their container
In other words, gases will spread out until they can’t spread out anymore
Gases will also move according to diffusion and effusion.

40. Racing Gases Demo Racing Gases Demo:
If concentrated HCl is at one end of the tube and concentrated NH3 is at the other end, which gas do you think will move farthest and fastest down the tube?
Racing Gases Demo
HCl (g)
NH3 (g)

41. RACING GASES DEMO The gases will diffuse down the tube
Diffusion – tendency of molecules to move from areas of higher concentration towards areas of lower concentration
Example: spraying perfume and smelling it across the room

42. DIFFUSION
Originally
Over Time

43. (Has a lower molar mass)
RACING GASES DEMO
The gases diffused at different rates
If the white ring forms closer to the HCl end of the tube, which gas moved farthest and fastest?
Why did this gas move faster?
BECAUSE IT’S LIGHTER!!
(Has a lower molar mass)

44. Graham’s Law of Effusion
The racing gases demo is related to Graham’s Law
Graham’s Law of Effusion – gases of lower molar masses effuse (and diffuse) faster than gases with higher molar masses
Effusion – when a gas escapes (diffuses) through a tiny hole in its container
Example: Helium balloons shrinking compared to normal balloons

45. The lighter the gas, the faster it moves!!!
Graham’s Law
Graham’s Law be applied to both the effusion and the diffusion of a gas
Gases with lower molar masses (lighter gases) diffuse faster than gases with higher molar masses (heavier gases)
The lighter the gas, the faster it moves!!!

46. Graham’s Law Which gas would diffuse and effuse faster…
Methane (CH4) or carbon dioxide (CO2)?
Chlorine (Cl2) or oxygen (O2)?
Hydrogen sulfide (H2S) or carbon monoxide (CO)?

47. Graham’s Law Formula
The rates are simply the speed or velocity at which the gas is traveling. So, this formula will compare the speed of one gas to the speed of the other gas.

48. Pressure of a Gas
The force of a gas exerted on the surface of a container
As gases bounce around in a container, each collision with the a container wall exerts a force
More collisions = higher pressure
Less collisions = lower pressure
Empty space with no particles and no pressure is called a vacuum

49. Pressure of a Gas
Atmospheric pressure – collision of air molecules with objects
As elevation increases, air density and therefore pressure decrease
Barometers measure atmospheric pressure

50. Pressure of a Gas
Vapor pressure – pressure due to force of gas particles above a liquid colliding with walls of container
Higher temp. = higher vapor pressure

51. Pressure of a Gas
When baking, there are different instructions for baking at high altitudes…why?
Boiling point of water decreases since lower pressure
Liquid leaves faster so food must bake longer
Gravity

52. Pressure of a Gas SI unit of pressure: pascal (Pa)
Other common pressure units:
Millimeters of mercury (mm Hg)
Atmospheres (atm)
1 atm = 760 mmHg = kPa = 760 torr

53. Practice Converting Units
1 atm = 760 mmHg = kPa
A tire pressure gauge records a pressure of 450 kPa. What is the pressure in atmospheres? In mm Hg?

54. DALTON’S LAW

55. (at constant volume and temperature)
Dalton’s Law
Partial pressure – the contribution of each gas in a mixture to the total pressure
Dalton’s Law of Partial Pressures – for a mixture of gases, the total pressure is the sum of the partial pressure of each gas in the mixture
Ptotal = P1 + P2 + P3 + …
(at constant volume and temperature)

56. Practice with Dalton’s Law
Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium. The partial pressures are: PO2= 20.0 kPa, PN2= 46.7 kPa, and PHe= 26.7 kPa.
Ptotal = P1 + P2 + P3 + …

57. Practice with Dalton’s Law
Air contains O2, N2, CO2, and trace amount of other gases. What is the partial pressure of oxygen (PO2) if the total pressure of the system is kPa and the partial pressures of N2, CO2, and the other gases are kPa, kPa, and 0.94 kPa, respectively?
Ptotal = P1 + P2 + P3 + …

58. Phase Changes What are phase changes?
A change in a substance’s state of matter
What are some examples of phase changes?

59. Phase Changes that Require Energy
If you have to put energy into a reaction to make it happen, it is called an endothermic reaction.
Endothermic Phase Changes:
Melting (solid  liquid) a.k.a “fusion”
Vaporization (liquid  gas)
Sublimation (solid  gas)

60. Phase Changes that Release Energy
If energy is released or given off by a reaction, it is called an exothermic reaction.
Exothermic Phase Changes:
Condensation (gas  liquid)
Deposition (gas  solid)
Freezing (liquid  solid)

62. Heating Curve