1. Chemical Bonding Notes

2. Neutral, “free” atoms are rarely found in nature
Most elements exist as part of a compound
Compounds are chemically bonded atoms

3. Chemical bond: strong attractive force that exists between atoms or ions in a compound
(+ nuclei and – electrons)

4. So, why do atoms bond?
To become more stable.
Bonding involves only valence electrons.

5. Remember…..
Noble gases are extremely STABLE because their outer shell is full of electrons
With the exception of He, all noble gases have 8 valence electrons

6. Octet Rule: Atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons

7. Types of Chemical Bonds
Ionic: formed when electrons are transferred from one atom to another
Usually formed between metals and non-metals.

8. Creates positive and negative ions which are electrostatically attracted

9. electrons are shared between atoms Usually formed between non-metals
2. Covalent: formed when pairs of
electrons are shared between atoms
Usually formed between
non-metals
Polar covalent: electrons are shared unequally
Non-polar covalent: electrons are shared equally

10. 3.Metallic: – results from attraction between positive ions and highly mobile, delocalized electrons
Delocalized electrons – valence electrons not held by any specific atom, but free to move from one atom to another
See “electron sea” model

11. Ionic & Covalent Properties
Ionic vs. Covalent___
“formula unit” “molecule”
high melting & low melting & boiling points boiling points
electrolyte non-electrolyte
hard & brittle pliable
Soluble in H2O polar: H2O soluble
non-polar: insoluble
in H2O

12. Metallic Properties
1. Luster: free electrons cause most metals to reflect light
2. High density: particles tightly packed in lattice
3. High melting and boiling points:
strong attractive forces between
positive ions and delocalized
electrons

13. 4. Good heat conductors: electrons transmit the energy vibrations
5. Good electrical conductors: mobile electrons flow in at one end and the same number flow out the other
6. Malleable and ductile: Can be bent into different shapes, and drawn into a long, thin wire. Distortion does not disrupt the bond

14. Ionic Bonds We use Electron Transfer Diagrams to represent Ionic Bonds
Ionization: The formation of an ion
1) Cation: positive ion formed by
the loss of one or more valence electrons
Metals tend to form + ions
2) Anion: negative ion formed by the
gain of one or more valence electrons
Non-metals tend to form - ions

15. Oxidation Numbers
The number of electrons that are transferred from an atom to form an ion.
This becomes the positive or negative charge of the ion.
Example:
Na  loses 1 valence electron
 becomes Na+1
It’s oxidation number is 1.

16. Electron Transfer Diagrams (used to represent ionic bonds)
To write:
1) use equation format
2) left of arrow: Show electron dot
diagrams
3) right of arrow: show ions
formed, and coefficients to give proper ion ratio

17. Examples
Na + Cl 
Mg + I 
Li + P 
Al + S 
Sr + O 

18. Lewis Structures (used to represent covalent bonds)
Formulas in which:
1)Element symbols represent the nuclei and core electrons.
2) Dot pairs (a dash) between two symbols represent covalent bonds (shared pairs)
3) Dot pairs around the outside of a symbol represent “lone pairs” (unshared pairs)

19. Lewis Structures (used to represent covalent bonds)
Single bond = 1 dash
Double bond = 2 dashes
Triple bond = 3 dashes
Central atom is the least electronegative atom.

20. Drawing Lewis Structures
Count valence electrons
Choose center atom
Connect atoms to central atom with single bond (dash)
Add unshared pairs around all surrounding atoms (except H).
Check for stability (each atom should have 8 valence electrons).
If not, rearrange unshared electrons creating multiple bonds where necessary.

21. Lewis Structure Examples:
HCl
NH3
CH3F
H2S
CO2

22. VSEPR Theory VSEPR: Valence Shell Electron Pair Repulsion
Theory states that the electrostatic repulsion between the valence-level electron pairs surrounding at atom causes these pairs to be oriented as far apart as possible.

23. VSEPR Theory
VSEPR allows for the prediction of geometric shapes of molecules.
“Lone Pairs” of electrons affect the shape but are not included in the shape.

24. Geometric Molecule Shapes
Linear
Bent
Trigonal pyramidal
Trigonal planar
Tetrahedral

25. Linear

26. Bent

27. Trigonal Pyramidal

28. Trigonal Planar

29. Tetrahedral

30. THE END!