1. Section 5.1—Types of Bonds

2. Ionic Bonding—Metal + Non-metal
Metals become cations (positively charged)
Non-metals become anions (negatively charged)
The cation & anion are attracted because of their charges—forming an ionic bond – ionic means formed from ions!!!

3. Covalent Bonding - between Non-metals
When two non-metals bond, they share electrons
Non-metals that share electrons evenly form non-polar covalent bonds
Non-metals that share electrons un-evenly form polar covalent bonds

4. Metallic Bonding – Metal to Metal Example: The Gold Atoms in a bar of Gold
The metals do not (as previously mentioned) hold onto their electrons well.
The valence electrons in metallic bonding are free to move throughout the structure - like a sea of electrons. This makes the atom appear to have a positive center.

6. Section 5.2—Drawing Molecules

7. Drawing Molecules on Paper
Lewis Structures (or Dot Structures) are one way we draw molecules on paper
A flaw of the Lewis method is that it is 2-D and molecules are 3D. However, Lewis Dots are essential to finding out the correct 3-D shape!

8. Step One: How many valence electrons are on each atom?
The main groups of the periodic table each have 1 more valence electron than the group before it. 1 2 3 4 5 6 7 8

11. Section 5.3—Molecules in 3D

12. Valence Shell Electron Pair Repulsion Theory (VSEPR Theory)
This theory (that bonds repel each other because they have like charges) attempts to explain why molecules form the shapes they form. Lewis Dot’s only allow us to see in 2-D! VSEPR allows us to visualize in 3-D!

13. What shapes do molecules form?
Linear
2 bonds, no lone pairs
Indicates a bond going away from you
Trigonal planar
3 bonds, no lone pairs
Indicates a bond coming out at you

14. What shapes do molecules form?
Tetrahedron
4 bonds, no lone pairs

15. Lone Pairs
Lone pairs are electrons, too…they must be taken into account when determining molecule shape since they repel the other bonds as well.
But only take into account lone pairs around the CENTRAL atom, not the outside atoms!

16. What shapes do molecules form?
Bent
2 bonds, 1 lone pair

18. What shapes do molecules form?
Bent
2 bonds, 2 lone pairs

19. What shapes do molecules form?
Trigonal pyramidal
3 bonds, 1 lone pair

20. Ionic Compound structures
Ionic compounds are made of positive and negative ions.
They pack together so that the like-charge repulsions are minimized while the opposite-charge attractions are enhanced.
Na+1
Cl-1

21. Section 5.4—Polarity of Molecules

22. Electronegativity
Definition: The pull an atom has for the electrons it shares with another atom in a bond.
Electronegativity is a periodic trend
As atomic radius increases and number of electron shells increases, the nucleus of an atom has less of a pull on its outermost electrons and the electrons within a bond!

23. Periodic Table with Electronegativies
increases
decreases

24. How to determine bond type
Find the electronegativies of the two atoms in the bond
Find the absolute value of the difference of their values
If the difference is 0.4 or less, it’s a non-polar covalent bond
If the difference is greater than 0.4 but less than 1.9, it’s a polar covalent bond
If the difference is greater than 1.9, it’s an ionic bond

25. Two atoms sharing equally
Each nitrogen atom has an electronegativity of 3.0 (from table)
They pull evenly on the shared electrons (You can tell this
by calculating the difference in their electronegativities. 3-3=0
so they share the electrons equally!
The electrons are not closer to one or the other of the atoms
This is a non-polar covalent bond

26. Atoms sharing almost equallyC H
Electronegativities: H = C = │ │= .4
The carbon pulls on the electrons slightly more, pulling them slightly towards the carbon
Put the difference isn’t enough to create a polar bond
This is a non-polar covalent bond

27. C O H Sharing unevenly Electronegativities: H = 2.1 C = 2.5 O = 3.5
The carbon-hydrogen difference isn’t great enough to create partial charges │ │= .4
But the oxygen atoms pulls significantly harder on the electrons than the carbon does. This does create a polar covalent bond │ │ = 1.0
This is a polar covalent bond

28. Showing Partial Charges
There are two ways to show the partial separation of charges
Use of “” for “partial”
Use of an arrow pointing towards the partial negative atom (THE ONE WITH GREATER ELECTRONEGATIVITY) and with a “plus” tail at the partial positive atom C O H + - C O H

29. Polar Bonds versus Polar Molecules
Not every molecule with a polar bond is polar itself
If the polar bonds cancel out then the molecule is overall non-polar.
The polar bonds cancel out.
No net dipole
The polar bonds do not cancel out.
Net dipole

30. The Importance of VSEPR
You must think about a molecule in 3-D (according to VSEPR theory) to determine if it is polar or not!
Water drawn this way shows all the polar bonds canceling out. O
H H
But water drawn in the correct VSEPR structure, bent, shows the polar bonds don’t cancel out!
Net dipole
H O H

31. Rules of Thumb for Molecular Polarity:
If the molecule is one of the basic VSEPR shapes (Linear, Trigonal Planar, Tetrahedron, trigonal bipyramid, and octahedral) and has all the SAME atoms connected to the central atom – the molecule is NONPOLAR
If the molecule has one of the basic VSEPR shapes and has DIFFERENT atoms then you should calculate the bond polarities to see if the molecule is polar.
The presence of a lone pair on the central atom ALMOST always make it POLAR unless the polarities cancel. (These exceptions will be addressed in AP Chem)

32. Section 5.5—Intermolecular Forces

34. Breaking Intramolecular forces
Breaking of intramolecular forces (within the molecule) is a chemical change
2 H2 + O2  2 H2O
Bonds are broken within the molecules and new bonds are formed to form new molecules

35. Breaking Intermolecular forces
Breaking of intermolecular forces (between separate molecules) is a physical change
Boiling water is breaking the intermolecular forces in liquid water to allow the molecules to separate and be individual gas molecules.

36. Dipole Forces
Polar molecules have permanent partial separation of charge.
The positive area of one polar molecule can be attracted to the negative area of another molecule. + - + -

38. Hydrogen Bonding – Strongest IMF
Why is it the strongest?
Hydrogen is a small atom and when its one electron attracts to one of 3 HIGH electronegative elements (F, O or N) that it is bonded to, its exposed nucleus can easily attract surrounding molecules!
NOTE: This ONLY happens when Hydrogen bonds with Nitrogen, Oxygen or Fluorine
(Remember as FON (phone )

39. Hydrogen Bond N H
Nitrogen – one of the elements required to have a Hydrogen bond
Intramolecular bond within the molecule
Hydrogen with “exposed” proton
Hydrogen bond –intermolecular
A lone pair of electrons that will attract strongly with the H on the OTHER molecule. N H

40. London Dispersion Forces
This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another.
Electrons move around the nuclei. They could momentarily all “gang up” on one side
All molecules have electrons. +
Positively charged nucleus -
Negatively charged electron + - + -
Electrons are fairly evenly dispersed. + -
As electrons move, they “gang up” on one side.

41. Strength of London Dispersion Forces
Electrons can gang-up and cause a non-polar molecule to be temporarily polar
The electrons will move again, returning the molecule back to non-polar
The polarity was temporary, therefore the molecule cannot always form LDF.
London Dispersion Forces are the weakest of the intermolecular forces because molecules can’t form it all the time. They are also called induced dipoles.

42. Strength of London Dispersion Forces
All molecules have electrons…all molecules can have London Dispersion Forces
The more electrons that gang-up, the larger the partial negative charge.
The larger the molecule, the stronger the London Dispersion Forces
Larger molecules have more electrons
Larger molecules have stronger London Dispersion Forces than smaller molecules.

43. IMF’s are Intermolecular Forces
Weakest - London Dispersion Forces
Dipole interactions
Strongest - Hydrogen bonding

44. Section 5.6—Intermolecular Forces & Properties

45. Bond type affects properties
The type of bonding affects the properties of the substance.
It takes energy to break and make bonds

46. IMF’s and Properties IMF’s are Intermolecular Forces
London Dispersion Forces
Dipole interactions
Hydrogen bonding
The number and strength of the intermolecular forces affect the properties of the substance.
It takes energy to break IMF’s

47. Melting/Boiling Points
Ionic bonds tend to have very high melting/boiling points as it’s hard to pull apart those electrostatic attractions
They’re found as solids under normal conditions

48. Melting/Boiling Points
Polar covalent bonds have the next highest melting/boiling points
Most are solids or liquids under normal conditions

49. Melting/Boiling Points
Non-polar covalent bonds have lower melting/boiling points
Most are found as liquids or gases

51. Water Water is a very small molecule
In general small molecules have low melting and boiling points
Based on its size, water should be a gas under normal conditions
However, because water is polar and can form dipole interactions and hydrogen bonding, its melting point is much higher
This is very important because we need liquid water to exist!

55. Solubility in Water
Ionic & polar covalent compounds tend to be soluble in water

57. Solubility in Water
Non-polar & metallic compounds tend to be insoluble

58. Solubility
In order for something to be dissolved, the solute and solvent must break the IMF’s they form within itself
They must then form new IMF’s with each other

59. Solubility - + - + - + - + + Solute, sugar (polar)
Solvent, water (polar) + -
- +
Solute, sugar (polar)
Water particles break some intermolecular forces with other water molecules (to allow them to spread out) and begin to form new ones with the sugar molecules.
- +
- +
- +

60. Solubility - + - + + Solute, sugar (polar) Solvent, water (polar)
- +
Solute, sugar (polar)
- +
As new IMF’s are formed, the solvent “carries off” the solute—this is “dissolving”

61. Solubility
If the energy needed to break old IMF’s is much greater than the energy released when the new ones are formed, the process won’t occur
An exception to this is if more energy is added somehow (such as heating)

62. Oil & Water Water is polar and can hydrogen bond, Oil is non-polar.
Water has London Dispersion, Dipole and hydrogen bonding. That takes a lot of energy to break
Water can only form London Dispersion with the oil. That doesn’t release much energy
Much more energy is required to break apart the water than is released when water and oil combine.
Therefore, oil and water don’t mix!

63. Conductivity of Electricity
In order to conduct electricity, charge must be able to move or flow
Metallic bonds have free-moving electrons—they can conduct electricity in solid and liquid state
Ionic bonds have free-floating ions when dissolved in water or when they are molten (liquid) form that allow them conduct electricity
Covalent bonds are NOT formed from charges and therefore cannot conduct electricity in any situation

64. Electrical Conductivity of solutions
No Light Dim Light Bright Light Bright Light
Question: Do all of these dissolve in water? If so, do all water soluble compounds conduct a current? Why or Why Not?

66. IMF’s and Viscosity Viscosity is the resistance to flow
Molasses is much more viscous than water
Larger molecules and molecules with high IMF’s become inter-twined and “stick” together more
The more the molecules “stick” together, the higher the viscosity

68. Surface Tension
Surface tension is the resistance of a liquid to spread out.
This is seen with water on a freshly
waxed car
The higher the IMF’s in the liquid,
the more the molecules “stick” together.
The more the molecules “stick” together,
the less they want to spread out.
The higher the IMF’s, the higher the surface tension.

70. Soap & Water Soap has a polar head with a non-polar tail
The polar portion can interact with water (polar) and the non-polar portion can interact with the dirt and grease (non-polar).
Polar head
Non-polar tail
Soap

71. Soap & Water
The soap surrounds the “dirt” and the outside of the Micelle can interact with the water.
The water now doesn’t “see” the non-polar dirt.
Dirt

75. Soap & Surface Tension
The soap disturbs the water molecules’ ability to form IMF’s and “stick” together.
This means that the surface tension of water is lower when soap is added.
The lower surface tension allows the water to spread over the dirty dishes.