1. Essential Questions for pages 302-312
1) What is electronegativity?
2) What is the group trend in it?
3) What is the period trend in it?
4) How is bond character determined by electronegativity?
5) How can you use it to determine bond type?
6) What are the 3 types on chemical bonds
7) Compare and contrast the 3 types of bonds.
8) What type(s) of bonding is involved with polyatomic ions.

2. Electronegativity (E.N.)
The attraction for an electron when you are bonded with another element.
Group trend↓: E.N. decreases as you go down due to size of atom
Period Trend→: E.N. increases as you go across due to increased nuclear charge
Differences in E.N. can determine the type of bond that holds the compound together.

3. Types of Chemical Bonds
Three types of bond:
Ionic bond: If E.N. difference is greater than eg: NaCl
Covalent bond: If E.N. difference is less than eg: CO2
Metallic bond: made of 2 metals and does not involve E.N.

4. Ionic Bonds Electronegativity difference  1.67.
Electrons are transferred from 1 element to another.
Tend to have high melting points.
Conduct electricity if melted or dissolved
Soluble in water
Crystals are sharply defined particles.
Usually made from a metal and a nonmetal.

5. Covalent Bonds electronegativity difference  1.67
Shares a pair or pairs of electrons to form molecules.
Properties include low melting points, does not conduct, and if solid it’s brittle
Usually made of 2 nonmetals or a nonmetal + Hydrogen
Carbon makes a lot of bonds C
Terms bond axis, bond length, and bond angle apply to them - bond is like a flexible spring attaching the atoms.

6. Metallic Bonds
Made from wandering electrons called delocalized electrons.
Always made of 2 metals.
Highly conductive, shiny, malleable.
Metal crystals form when metals crowd and outer electrons overlap - called delocalized electrons. This causes luster (shimery) and malleability.
Alloy: combination of 2 or more metals.
It strengthens metals with few delocalized electrons.
Examples are steel, bronze, brass
Requires heat to speed the mixing of the electrons on the surface.
Alloys always have different properties from the pure metals.

7. Polyatomic Ions
Polyatomic ions are covalently bonded but then the molecule gets charges by gained extra electrons.

8. Distance between nucleus and valence electrons.
Particle Size – comparing 4 types of radii
Atomic Radii
Ionic Radii
Covalent Radii
Van der Waals Radii
Distance between nucleus and valence electrons.
It’s measured on atoms in crystals.
Trend was discussed in Cpt. 10
Differs from atomic radii because atoms gain or loose electrons.
Cl- is bigger than Cl
Ca is bigger than Ca 2+
To find internuclear distance between 2 ions in an ionic compound: add the 2 ionic radii in the compound.
To find ionic radii: take ½ of the internuclear distance.
Only approx since e- are fuzzy and ions interact
Radii of atoms in a molecule.
In a molecule, covalent radii is ½ internuclear distance
To find bond length: it’s the sum of covalent radii from table 12.5.
Expect covalent radii to be less than atomic radii, but cloud can distort and be bigger on occasion.
Radii of unbound atoms.
It’s the minimum distance between atoms since e- repel.
Like an imaginary, rigid shell around an unbound atom, at the point where electron probability drops to lower than 90%.
Expect Van der Walls radii to be less than atomic radii.
Do page #9-13

9. Unlike a ball, an atom doesn't have a fixed radius
Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.
                                                                           
As you can see from the diagrams, the same atom could be found to have a different radius depending on what was around it.
The left hand diagram shows bonded atoms. The atoms are pulled closely together and so the measured radius is less than if they are just touching. This is what you would get if you had metal atoms in a metallic structure, or atoms covalently bonded to each other. The type of atomic radius being measured here is called the metallic radius or the covalent radius depending on the bonding.
The right hand diagram shows what happens if the atoms are just touching. The attractive forces are much less, and the atoms are essentially "unsquashed". This measure of atomic radius is called the van der Waals radius after the weak attractions present in this situation.

10. Cpt. 13 Molecular Structure
Objective: To find the shapes of molecules
1) Electron Distribution
A) Use Lewis electron dot diagrams to “see” electron arrangement to predict shape
B) For all atoms, except H and He, 8 electrons is a full outer level.
C) electrons in the dot diagram are “shared pairs” (in between 2 atoms and forms the covalent bond) or “unshared pairs” (not attached to a second atom)

11. Example
NH3
Do page 323 #1-2.

12. 2) Electron Pair Repulsion (VSEPR Theory)
A) Electron pairs spread as far apart as possible since like charges repel.
All repulsion is not equal: unshared pairs have the most repulsion because they form the largest electron cloud.
Order of repulsion from greatest to least:
Unshared–unshared unshared-shared  shared-shared

13. Example of VSEPR Theory
Example of CH4 (all shared) versus NH3

14. 3) Molecular Shapes A) Linear
Contains 2 shared pairs and 0 unshared pairs (or just 1 shared pair.)
Bond angle is 180 degrees
Eg: BeF2 or H2
B) Trigonal Planar
Contains 3 shared pairs and 0 unshared pairs
Bond angle is 120 degrees
eg: GaF3
C) Bent
Contains 2 shared pairs and 1 unshared pairs
Bond angle is 120
eg: SiH2 and O3 (ozone)

15. Shapes, Cont’s D) Tetrahedral
Contains 4 shared pairs and 0 unshared pairs
Bond angle is 109.5
eg: CH4 (like a pyramid)
E) Trigonal Pyramidal
Contains 3 shared pairs and 1 unshared pairs
Bond angle is 109.5
eg: NH3 (pyramid with distorted top)
F) Bent (version 2)
Contains 2 shared pairs and 2 unshared pairs
eg: H2O (pyramid with distorted top and 1 base corner)
Do page 328 #3.

16. Geometry of Carbon Compounds
Carbon makes “hybrid orbitals”. One electron leaves its orbital to form another p orbital. s2p2 becomes the hybrid s1p3.
Dot Diagram of the carbon hybrid. ●
● C ●
Other atoms form hybrids also, any time its needed to bond onto an available atom in a covalent bond.

18. Double and Triple Bonding
Carbon can also form double and triple bonds.
Single bond shares 1 pair of electrons.
Double bond shares 2 pairs of electrons.
Triple bonds share 3 pairs of electrons.
The more pairs shared, the stronger the bond and the shorter the bond length. The molecular shape does not change drastically, but can bend some since the double or triple bond electron cloud is larger.

19. Compounds that have only single bonds are called “saturated” (like a saturated fat).
Compounds that have at least 1 double or triple bond are called “unsaturated”.
Do page 334 # 4-6.
Do handout, Geometry of methane molecule.

20. Other Forces at Work in Molecules:
Intramolecular Forces - hold atoms in molecules (internal bonding – Ionic, covalent, and metallic bonds)
Intermolecular Forces - hold molecules to each other. (external)

21. An Intermolecular Force : Polar or Nonpolar
Polar - means opposite charged ends.
Some molecules are polar (like H2O)
Polar bonds form if there are big differences in electronegativity - 1 side hogs the electrons making it slightly more negative than the other side. The opposite side then is slightly more positive.
e.g H2O H O H

22. Polar, cont’d The bond is then called polar covalent.
A polar molecule is also called a dipole.
You can show the partial charges by Greek symbol delta H O H

23. Polar, still Problem: Is I2 a polar or nonpolar bond?
Look up electronegativities and subtract
= 0 so its nonpolar
Do pg 352 #1 and 2 (Use page 303)
You can have a polar bond but it may or may not form a polar molecule - polar molecules are not symmetrical. (look at dot diagram)
Symmetrical shaped molecules will cancel out the dipoles formed from your electronegativity differences

25. Rule of thumb - “Like dissolves like.”
A polar molecule can only dissolve other polar molecules (& also ionic bonds)
An Intermolecular Force - van der Waals forces: Weak forces caused by dipole-dipole interaction.

26. This is why “like dissolves like”+ _ - + + - _ + + - + _

27. Applications/Other Terms
Miscibility – the ability of two or more substances to mix, and form a single homogeneous phase. When two substances are immiscible they will form separate phases when mixed; the best known example is oil and water.
Chromatography: Only works when “likes dissolves like”. A polar solvent will only work to separate a polar substance on the substrate (the paper, gas, liquid used to separate the substance)