1. bonding

2. Electronegativity- the ability of an atom to attract a shared pair of electrons to itself. The difference in electro negativities of 2 elements can determine the type of bond that holds them together.

3. Differences in electronegativity
= covalent bond(nonpolar)
.41 – 1.66 = polar covalent bond
1.67 or greater = ionic bond

4. Chemical bonds
The force of attraction between the electrons of one atom and the nucleus of another atom.
3 types
Ionic
Covalent
metallic

5. Bond Types Sigma bonds () – single covalent bond
Pi bonds () – occur when multiple bonds are formed
Single bond – sigma
Double bond – 1 sigma & 1 pi
Triple bond – 1 sigma & 2 pi

6. Pi Bonding in the Nitrate Ion
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7. Pi Bonding and Antibonding Orbitals
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8. Covalent Bonding Covalent bond –formed from the sharing of electrons
Usually between 2 non metallic elements

9. Covalent bonds When 2 nonmetals combine.
Both want the electrons available.
The electrons in this type of bond are shared so that both elements can feel “full”
These types of bonds produce molecules.

10. polarity
The ability of an atom in a covalent bond to pull the shared pair of electrons more strongly to itself.

11. Molecules Molecule – two or more atoms covalently bound together
Diatomic molecule – two of the same atom bound together H, N, O, F, Cl, Br, I
H2, N2, O2, F2, Cl2, Br2, I2

12. Octet & Duet Rules Octet Rule – atoms want to have 8 valence electrons
Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

13. Octet rule exceptions
3. B prefers 6 e’.

14. Ionic bonds When a metal and a nonmetal combine.
E’ from the metal are transferred to the nonmetal.
The goal is for both elements to achieve the noble gas configuration.

15. 1. Predict if the following bonds are ionic or covalent
Al-Si Li-S Ca-Cl
Ba-O Ca-P F-S
C-H B-Na Br-Rb

16. Comparison of a Molecular Compound and an Ionic Compound
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17. Metallic bonds
The electrons in this type of bond to not center around one atom but appear to float around each nucleus as a delocalized cloud. This causes metallic bonds to exhibit unique properties.
1. malleability
2.ductility

18. 3. good conductor of heat and electricity.
4. hardness is determined by bond strength.
5. luster

20. Bond length & Strength
As the number of bonds increases, the bond length decreases
The shorter the bond, the stronger the bond
Single> double>triple

22. Lewis diagrams & the octet rule
Gilbert Newton Lewis (surrounded by pairs of electrons)

23. Lewis Structure
Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule
There are rules for Lewis Structures that are based on the formation of a stable compound
Atoms want to achieve a noble gas configuration

24. Steps for drawing Lewis Structures
Sketch a simple structure with a central atom and all attached atoms
Add up all of the valence electrons for each individual atom
If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge
If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

25. Steps for drawing Lewis Structures
Subtract 2 electrons for each bond drawn
Complete the octet on the central atom & subtract those electrons
Complete the octet on the surrounding atoms & subtract those electrons
Get your final number
If 0  you are done!
If +  add that many electrons to the central atom
If -  need to form multiple bonds to take away that many electrons

26. Examples
CCl4
Sketch a simple structure with a central atom and all attached atoms Cl │
Cl – C – Cl

27. Examples Add up all of the valence electrons for each individual atom
4 + 4(7) = 32
Subtract 2 electrons for each bond drawn
32-8 = 24
Complete the octet on the central atom & subtract those electrons
Done

28. Examples
Complete the octet on the surrounding atoms & subtract those electrons
24 – 24 = 0
Final number = 0…DONE!
Final structure is… __
│Cl │
__ │ __
│Cl – C – Cl │ │

29. Draw the Lewis structures for the following molecules:
HCl
SCl2
PH3
SiF4

30. Polyatomic ion Lewis structures
1. determine the total number of valence electrons involved.
2. add the charge of the polyatomic to the total number.
3. arrange the elements as they would bond together.
4. place a pair of electrons everywhere a bond occurs.

31. Use the remaining electrons to fill the octet rule for the elements involved in the bond.
* remember, electrons always travel in pairs.

32. When there are not enough e’ to satisfy everyone’s octet, double and triple bonds are formed!

33. Formation of C=C Double Bond in Ethylene
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34. 2. Draw the Lewis dot diagrams for the elements Z=7 thru Z=20
3. Draw the Lewis dot diagrams for each of the following polyatomic ions
NH4+ OH- C2O O2-2

35. Molecular formulas- Used to tell exactly how many of each atom is present in the compound.
Empirical Formulas- The simplest formulas possible indicating the ratio of atoms found in the molecule. Think of it as reducing a fraction to the common denominator.

36. Write the empirical formula for the following:
C6H6
C2H2
C6H12O6
C4H10
P4O10
SO3
N2O4
NO2
Ag2C4H4O6
K2S4