1. Physical State of Matter
Honors Chem

2. States of Matter

3. States of Matter
The fundamental difference between states of matter is the strength of the intermolecular forces of attraction.
Stronger forces bring molecules closer together.
Solids and liquids are referred to as the condensed phases.

4. Differences in the States of Matter

5. Physical States: A Molecular Model

6. Solids: Crystalline vs. Amorphous
Amorphous -  Amorphous solids have irregular or curved surfaces, do not give well-resolved x-ray diffraction patterns, and melt over a wide range of temperatures.
Crystalline – Crystalline solids have well-defined edges and faces, diffract x-rays, and tend to have sharp melting points.

7. Effects of Temperature and Pressure
By heating and cooling a substance you can change its state - as temperature changes directly affect the average kinetic energy.
Increasing pressure can change the state of a substance
higher pressure will bring molecules closer together, making their intermolecular interactions more effective.
At 1 atm (25oC) propane is a gas but at higher pressures (25oC) propane is found in a liquid state.

8. Which State of Matter?
The answer to this question largely relies on the
balance between the kinetic energies of the particles.
interparticle energies of attraction.

9. Some Electronegativity Values

10. Molecular Geometry and Electron-Pair Geometry

11. Definitions: Molecular Geometry: the shape of the molecule
to determine: write the Lewis structure and determine the number of bonding groups of electrons and the number of non-bonding pairs of electrons on the central atom, then use the associated name for that shape.
Electron-Pair Geometry: provides a guide to the bond angles of a terminal-central-terminal atom in a compound
* there are several examples with the same electron-pair geometry, but different molecular geometries

12. Why is this Important?
The study of molecular geometry is important in that a molecule’s geometry affects its physical and chemical properties, such as melting point, boiling point, density, and the types of reactions it undergoes

13. Summary of molecular / electron-pair geometries for different combinations of bonding groups and nonbonding pairs of electrons on the central atom.
# of bonding groups/domains on 'central' atom
# of lone pair electrons on 'central' atom
Electron-pair Geometry
Molecular Geometry
Bond
Angle 2
Linear
180 3
Trigonal planar
120 1
Bent
Less than 120 4
Tetrahedral
109.5
Trigonal pyramidal
Less than 109.5

14. Definitions:
Lone pair: an electron pair in the outermost shell of an atom that is not shared or bonded to another atom.
Bonding Pair: pairs of electrons that are shared between atoms in a molecule and hold the atoms together
Valence Electrons: outermost electrons of an atom, involved in forming bonds and making molecules.
VSEPR (valence-shell electron-pair repulsion) model: used to determine a molecule's general geometry, used to predict the way electrons within atoms will repel each other

15. Two Electron Groups In a molecule of BeCl2,
2 bonds, 0 lone pairs
In a molecule of BeCl2,
there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule).
To minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other. The shape of the molecule is linear.

16. Three Electron Groups In a molecule of BF3,
3 bonds, 0 lone pairs
In a molecule of BF3,
three electron groups are bonded to the central atom B (B is an exception to the octet rule).
repulsion is minimized with 3 electron groups at angles of 120°.
the shape is trigonal planar.

17. Three Electron Groups with a Lone Pair
In a molecule of SO2,
S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair.
repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement.
the shape is determined by the two O atoms bonded to S, giving SO2 a bent (~120°) shape.

18. Four Electron Groups In a molecule of CH4,
4 bonds, 0 lone pairs
In a molecule of CH4,
there are four electron groups around C.
repulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangement.
the four bonded atoms form a tetrahedral shape.

19. Four Electron Groups with a Lone Pair
In a molecule of NH3,
three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair.
repulsion is minimized with 4 electron groups in a tetrahedral arrangement.
the three bonded atoms form a pyramidal (~109°) shape.

20. Four Electron Groups with Two Lone Pairs
In a molecule of H2O,
two electron groups are bonded to H atoms and two are lone pairs (4 electron groups).
four electron groups minimize repulsion in a tetrahedral arrangement.
the shape with two bonded atoms is bent (~109°).

21. Polar vs. Nonpolar

22. Polar Molecules A polar molecule contains polar bonds.
has a separation of positive and negative charge. called a dipole, indicated with δ+ and δ–.
has dipoles that do not cancel.
δ+ δ–
H–Cl NH3
dipole
Dipoles do not cancel.

23. Nonpolar Molecules A nonpolar molecule contains nonpolar bonds
Cl–Cl H–H
or has a symmetrical arrangement of polar bonds.

24. Predicting Molecular Polarity

25. Step 1: Draw a Lewis structure
Step 2: Identify each bond as either polar or nonpolar.
If there are no polar bonds, the molecule is nonpolar.
If the molecule has polar bonds, move on to Step 3.
Step 3: If there is only one central atom, examine the electron groups around it.
If there are no lone pairs on the central atom, and if all the bonds to the central atom are the same, the molecule is nonpolar.
If the central atom has at least one polar bond and if the groups bonded to the central atom are not all identical, the molecule is probably polar. Move on to Step 4.
Step 4: Draw a geometric sketch of the molecule.
Step 5: Determine the symmetry of the molecule using the following steps.
Describe the polar bonds with arrows pointing toward the more electronegative element. Use the length of the arrow to show the relative polarities of the different bonds. (A greater difference in electronegativity suggests a more polar bond, which is described with a longer arrow.)
Decide whether the arrangement of arrows is symmetrical or asymmetrical
If the arrangement is symmetrical and the arrows are of equal length, the molecule is nonpolar.
If the arrows are of different lengths, and if they do not balance each other, the molecule is polar.
If the arrangement is asymmetrical, the molecule is polar
If difference in electronegativity for the atoms in a bond:
>0.4 bond = polar.
<0.4 bond = nonpolar

26. Ex) Molecular Polarity of H2O
Step 1: Lewis Structure

27. Ex) Molecular Polarity of H2O
Step 2: Identify each bond as either polar or nonpolar
From the electronegativity table:
O = H = 2.1
3.5 (O) (H) = 1.4
which makes the O — H bonds, polar covalent.

28. Ex) Molecular Polarity, H2O
* A molecule can possess polar bonds and still be nonpolar. If the polar bonds are evenly (or symmetrically) distributed, the bond dipoles cancel and do not create a molecular dipole
Ex) Molecular Polarity, H2O
Step 5: Describe the polar bonds with arrows pointing toward the more electronegative element If the bonds are polar covalent, draw the electron-dot formula and determine if the dipoles cancel or not. The four electron groups of oxygen are bonded to two H atoms. Thus, the H2O molecule has a net dipole, which makes it a polar molecule.

29. Learning Check
Determine the shape of each of the following molecules and whether they are polar or nonpolar. Explain.
1. PBr3
2. HBr
3. Br2
4. SiBr4

30. Solution
Determine the shape of each of the following molecules and whether they are polar or nonpolar. Explain.
1. PBr3 pyramidal; polar; dipoles don’t cancel
2. HBr linear; polar; one polar bond (dipole)
3. Br2 linear; nonpolar; nonpolar bond
4. SiBr4 tetrahedral; nonpolar; dipoles cancel

31. Intermolecular Forces

32. Inter vs. Intramolecular Forces
Intramolecular forces: those forces between the atoms found inside a single molecule.
Example: Metallic, ionic and covalent bonds
Intermolecular forces: those forces between molecules.
These forces determine the boiling point of substances thus the state.
Examples: dipole­ dipole forces, London dispersion forces, hydrogen bonding

33. Intermolecular Forces
The attractions between molecules are not nearly as strong as the intramolecular attractions (bonds) that hold compounds together.
Many physical properties reflect intermolecular forces, like boiling points, melting points, viscosity, surface tension, and capillary action.

34. Intermolecular Forces
Less energy is required to vaporize a liquid or melt a solid than to break the individual bonds.
To vaporize HCl requires 16 kJ/mol
To break the covalent bond in HCl requires 431 kJ/mol

35. Types of Intermolecular Force
Weakest to strongest forces:
dispersion forces (or London dispersion forces)
dipole–dipole forces
hydrogen bonding (a special dipole–dipole force)
ion–dipole forces
Note: The first two types are also referred to collectively as van der Waals forces.

36. Dispersion Forces
The figure below shows how a nonpolar particle (in this case a helium atom) can be temporarily polarized to allow dispersion force to form.
The tendency of an electron cloud to distort is called its polarizability.

37. Factors Which Affect Amount of Dispersion Force in a Molecule
number of electrons in an atom (more electrons, more dispersion force)
size of atom or molecule/molecular weight
shape of molecules with similar masses (more compact, less dispersion force)

38. Polarizability & Boiling Point
If something is easier to polarize, it has a higher boiling point.
Remember: This means more intermolecular force (larger molecule: larger molecular weight, more electrons).

39. Dipole–Dipole Interactions
*only in polar molecules
Polar molecules have a more positive and a more negative end–a dipole (two poles, δ+ and δ−).
The oppositely charged ends attract each other.

40. Dipole–Dipole Interactions
For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point.

41. Which Have a Greater Effect: Dipole–Dipole Interactions or Dispersion Forces?
Dispersion forces are found in all substances and make up a larger contribution to the total intermolecular attractions
If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force.
If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

42. Hydrogen Bonding
The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong.
We call these interactions hydrogen bonds.
A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

43. What Forms Hydrogen Bonds?
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
These atoms interact with a nearly bare nucleus (which contains one proton).

44. Hydrogen Bonds
Hydrogen bonds are much weaker than covalent bonds but stronger than dispersion forces and dipole-dipole forces.
This leads to their huge biological significance
Hydrogen bonding is responsible for stabilizing structures of proteins and for the way DNA is able to carry genetic information

45. Hydrogen Bonding

46. Checkpoint
In which of these molecules is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl fluoride (H3CF), and hydrogen sulfide (H2S)?

47. Answer to Checkpoint:
H2NNH2, because it has Hydrogen bonds (H is directly bound to N)
NOT:
CH4 because C does not count as one of the atoms H can bind to to form a Hydrogen bond
H2S because S does not count as one of the atoms H can bind to to form a Hydrogen bond
H3CF because there is no H-F bond

48. Ion–Dipole Interactions
Ion–dipole interactions are found in solutions of ions.
The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

49. Summarizing Intermolecular Forces

50. Checkpoint
List the following substances in order of increasing boiling point: BaCl2, H2, CO, HF, and Ne.

51. Answer to Checkpoint
H2 (dispersion) < Ne (dispersion) < CO (dipole-dipole) < HF (Hydrogen bonds) < BaCl2 (Ionic)
Because boiling point increases as the strength of the intermolecular forces increases
Ne has a stronger dispersion force than H2 because it has more electrons (refer to slide 37 for more info)
CO is polar because O is more electronegative than C

52. Liquid Properties and Intermolecular Forces

53. Liquid Properties Affected by Intermolecular Forces
boiling point (previously discussed) and melting point
viscosity
surface tension
capillary action

54. Viscosity Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

55. Viscosity
The shapes and flexibility of molecules play a role as well as they tend to become entangled.
For a series of related compounds viscosity increases with molecular weight.
The SI unit for viscosity is kg/m-s

56. Surface Tension
Water acts as if it has a “skin” on it due to extra inward forces on its surface. Those forces are called the surface tension.

57. Surface Tension
Molecules in the interior are attracted equally in all directions whereas molecules on the surface experience a net inward force
This net force tends to pull surface molecules towards the interior , thereby reducing the surface area and making the molecules at the surface pack closer together

58. Surface Tension
Surface tension is the energy required to increase the surface area of a liquid by a unit amount
The surface tension of water at 20oC is 7.29 x 10-2 J/m2.
Water has a high surface tension due to its hydrogen bonds

59. Cohesion and Adhesion – Capillary Action
Intermolecular forces that bind similar molecules to one another are called cohesive forces.
Intermolecular forces that bind a substance to a surface are called adhesive forces.
These forces are important in capillary action.

60. Capillary Action
The rise of liquids up narrow tubes is called capillary action.
Adhesive forces attract the liquid to the wall of the tube.
Cohesive forces attract the liquid to itself.
Water has stronger adhesive forces with glass; mercury has stronger cohesive forces with itself.

61. Phase Changes

62. Phase Changes
Conversion from one state of matter to another is called a phase change.
Energy is either added or released in a phase change.
Phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.
Start here 6/21/10

63. Phase Change
In a solid, particles are in a fixed position and closely arrange to minimize the energy of the system
As the temperature increases, the particles vibrate with increasing energetic motion
When the solid melts, the particles are free to mole relative to one another, increasing their average kinetic energy

64. Energy Change & Change of State
The heat of fusion is the energy required to change a solid at its melting point to a liquid.
The heat of vaporization is the energy required to change a liquid at its boiling point to a gas.
The heat of sublimation is the energy required to change a solid directly to a gas.

65. Sources http://preparatorychemistry.com/Bishop_molecular_polarity.htm