1. Chapters 8.4, 15.1, and 13.3 “States of Matter”

2. Section 8.3 The Condensed States of Matter
OBJECTIVES:
Explain how the Kinetic Molecular Theory accounts for the physical properties of Liquids and Solids.

3. Section 8.3 The Condensed States of Matter
OBJECTIVES:
Describe the different types of Intermolecular Forces

4. Condensed States of Matter
In the previous chapter, you studied the physical behavior of gases.
Now you are going to learn about liquids and solids.
Liquids and solids are referred to as the condensed states of matter because substances in these states have substantially higher densities than they do in the gaseous states.

5. GAS LIQUID SOLID Highly compressible Slightly compressible Low density
High density
Fills container completely
Definite Volume
Rigidly retains its volume
Takes the shape of its container
Retains its own shape
Rapid Diffusion
Slow Diffusion
Extremely slow Diffusion
High expansion on heating
Low expansion on heating

6. Intermolecular Forces
IMF are forces of attraction that exist between neighboring molecules in a compound.
Remember….
Ionic Bonds -solids(positive and negative ions)
Metallic Bonds -solids(neg. electrons and protons)
Covalent Bonds -solids, liquids, and gases
These three types of bonding are INTRAMOLECULAR Forces

7. Ionic Compounds Electrons from metals to non-metals.
Lose 1 or more electrons to achieve noble gas configuration.
This transfer is called an IONIC BOND.
All Ionic Compounds are solids at room temperature. So what does this tell you about the strength of attraction?
It is very STRONG.

8. Metallic Bonds Holds atoms together in a metallic substance.
Share valence electrons among themselves.
Are these bonds strong?
Yes, just about-except Mercury (liquid)

9. Covalent Bonding Holds Gases and Liquids together.
Share electrons to achieve noble gas configuration.
Contains Intermolecular forces.
The force of attraction between neighboring molecules is Intermolecular forces.

10. Intermolecular Forces
There are three types of Intermolecular Forces:
Dispersion Forces
Dipole-Dipole Forces
Hydrogen Bonding

11. Dispersion Forces
A Dispersion Force is an attraction between induced dipoles.
A Dipole consists of a separation of the two opposite charges by some distance.
An Induced Dipole is a dipole that is created by the presence of a neighboring dipole.

13. Dispersion Forces
A symmetrical, spherical atom has no separation of charge.
The electrons are distributed uniformly around the nucleus. At any time, however, such an atom can lose its spherical shape and become a temporary dipole.
A temporary dipole can distort the electron cloud of a symmetrical atom and induce a dipole in it.

14. Dispersion Forces Examples: Diatomics, Non-Polar
The higher the mass, the higher the dispersion force. (Br2 > H2)
The larger the atom, the higher the boiling point needed to break the dispersion force.

15. Dipole-Dipole Forces
Dipole-Dipole Forces are attractions between opposite charges of neighboring permanent dipoles.
A permanent dipole is the result of a covalent bond between 2 different atoms.
Dispersion forces depend on temporary dipoles induced in neighboring atoms.

16. Dipole-Dipole Forces Molecules with permanent dipoles are polar.
A permanent dipole is a result of a covalent bond between two different atoms.
Because 2 atoms are not identical, there is a difference in their electronegativities, so…. Each atom acquires a partial charge.

19. Dipole-Dipole Forces
The greater the difference in electronegativities, the greater the dipole.
All Diatomic elements that are included with polar bonds constitute permanent dipoles.
Examples: Cl2O, NO or anything polar
These bonds are strong and require a higher boiling point than dispersion forces to break them.

20. Hydrogen Bonding Any bond with Hydrogen.
Hydrogen with any Halogen has a large polarity because of an electronegativity difference.
Such bonds with Hydrogen, have higher boiling points than can be accounted for either by dispersion forces or by dipole-dipole forces.

21. Hydrogen Bonding
Since there is such a large electronegativity difference, there is a stronger force involved…Hydrogen Bonding.
Not a Covalent Bond, but a Strong Intermolecular Force.
Example: water, HF, NH3
Hydrogen Bonding is the strongest of all the forces.

23. 15.1 Properties of Liquids
The physical properties of Liquids are determined mainly by the nature and strength of the intermolecular forces present between their molecules.
Some factors include:
Viscosity
Surface Tension

24. Viscosity
The “friction,” or resistance to motion, that exists between molecules of a liquid when they move past each other.
Molasses High viscosity
Gasoline Low viscosity

26. Viscosity Liquids with hydrogen bonding have higher viscosities.
Gasoline molecules are larger and have dispersion forces.
Viscosity increases as temperature decreases. (Relates back to KMT)

27. Surface Tension
Molecules at the surface of a liquid experience attractive forces downward, toward the inside of the liquid, and sideways, along the surface of the liquid.
The imbalance of forces at the surface of a liquid results in a property called surface tension.

28. Surface Tension
The uneven forces makes the surface behave as If it had a tight film stretched across it.
Razor blades, needles, small bugs can float on water.
Surface tension is greater in liquids with a strong intermolecular force of attraction.
Surface tension increases as temperature decreases.

30. Unusual Properties of Water
Water has a higher boiling point, so it is a liquid at room temperature.
Water can absorb or release relatively large quantities of heat without large changes in temperature.
The density of the solid form of water, ice, is less than the density of its liquid form, water.

31. Unusual Properties of Water
Ice is also a good insulator. (more open structure in air)
Water has a relatively high surface tension.
Water has a very high heat of vaporization. (The amount of heat required to convert a given amount of a liquid into a gas. This is due to strong IMF)
A Universal Solvent because of its polarity.

32. Section 13.3 The Nature of Solids
OBJECTIVES:
Evaluate how the way particles are organized explains the properties of solids.

33. Section 13.3 The Nature of Solids
OBJECTIVES:
Identify the factors that determine the shape of a crystal.

34. Section 13.3 The Nature of Solids
OBJECTIVES:
Explain how allotropes of an element are different.

35. Particles in a liquid are relatively free to move.
The Nature of Solids
Particles in a liquid are relatively free to move.
Solid particles are not.
Solid particles tend to vibrate about fixed points, rather than sliding from place to place.

36. The Nature of Solids
Most solids have particles packed against one another in a highly organized pattern
Tend to be dense and incompressible
Do not flow, nor take the shape of their container
Are still able to move, unless they would reach absolute zero

37. The Nature of Solids
When a solid is heated, the particles vibrate more rapidly as the kinetic energy increases.
The organization of particles within the solid breaks down, and eventually the solid melts.
The melting point (mp) is the temperature a solid turns to liquid.

38. The Nature of Solids
Generally, most ionic solids have high melting points, due to the relatively strong forces holding them together
Sodium chloride (an ionic compound) has a melting point = 801 oC
Molecular compounds have relatively low melting points

39. 2 Main Types of Solids
Crystalline Solids
Amorphous Solids

40. Amorphous Solids
Some substances are rigid and appear solid but do not behave like crystalline solids.
The word amorphous comes from the greek word meaning “without form.”
Glass, rubber, some plastics
Liquids that have been cooled…glass.

42. Crystalline Solids
A solid in which the representative particles exist in a highly ordered, repeating pattern.
Most solid substances are crystalline in nature.
Snow, sugar, salt, precious stones, almost all of the metals

43. Crystalline Solids
In a crystal, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three-dimensional pattern called a crystal lattice
All crystals have a regular shape, which reflects their arrangement

44. Crystalline Solids
The type of bonding that exists between the atoms determines the melting points of crystals
A crystal has sides, or faces
The angles of the faces are a characteristic of that substance, and are always the same for a given sample of that substance

45. Crystalline Solids Crystals are classified into seven groups.
The 7 crystal systems differ in terms of the angles between the faces, and in the number of edges of equal length on each face.

46. Crystalline Solids
The shape of a crystal depends upon the arrangement of the particles within it
The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell

48. Crystalline Solids
There are three kinds of unit cells that can make up a cubic crystal system:
1. Simple cubic
2. Body-centered cubic
3. Face-centered cubic

49. - Page 398

50. Crystalline Solids
Some solid substances can exist in more than one form
1. Diamond, formed by great pressure
2. Graphite, which is in your pencil
3. Buckminsterfullerene (also called “buckyballs”) arranged in hollow cages like a soccer ball

51. These are called allotropes of carbon, because all are made of pure carbon only , and all are solid.
Allotropes are two or more different molecular forms of the same element in the same physical state.

52. Section 14.4 Changes of State
OBJECTIVES:
Identify the conditions necessary for sublimation.

53. Section 14.4 Changes of State
OBJECTIVES:
Describe how equilibrium conditions are represented in a phase diagram.

54. End of Chapter 14