1. METALS Shiny 'metallic' appearance
Solids at room temperature (except mercury)
High melting points
High densities
Large atomic radii
Low ionization energies
Low electronegativities
Usually, high deformation
Malleable
Ductile
Thermal conductors
Electrical conductors

2. Non-Metals No luster (dull appearance)
Poor conductor of heat and electricity
Brittle (breaks easily)
Not ductile
Not malleable
High ionization energies
Low density
Low melting point
High electronegativities

3. Semi-metals (Metalloids)
Solids (at room temperature)
Can be shiny or dull
Ductile
Malleable
Conduct heat and electricity better than nonmetals but not as well as metals

4. Electronegativity
The electronegativity of an atom is a measure of its power when in chemical combination to attract electrons to itself.
With few exceptions, electronegativity increases from left to right across the periodic table and decreases down a group:
F is far more electronegative than I
F is far more electronegative than Li

5. Broad Classification of Chemical Bonds:
Ionic bonds - electrostatic forces that exist between ions of opposite charge
typically involves a metal with a nonmetal. Example: NaCl.
Covalent bonds - results from the sharing of electrons between two atoms typically involves one nonmetallic element with another. Example: Cl2.
Metallic bonds -each metal atom bonded to several neighboring atoms so that bonding electrons are free to move throughout the three-dimensional structure. Example: Cu.

6. Lewis and Langmuir
Two Americans who were instrumental in developing a bonding theory based on the number of electrons in the outermost "valence" shell of the atom were Gilbert Newton Lewis (1875–1946) and Irving Langmuir (1881–1957).

7. In 1902, while Lewis was trying to explain valence to his students, he depicted atoms as constructed of a concentric series of cubes with electrons at each corner. This "cubic atom" explained the eight groups in the periodic table and represented his theory that chemical bonds are formed by electron transference to give each atom a complete set of eight.
G. N. Lewis at his desk

8. Langmuir elaborated on ideas first expressed by Lewis
Langmuir elaborated on ideas first expressed by Lewis. Langmuir proposed that octets could be filled by sharing pairs between two atoms—the "covalent" bond.
Irving Langmuir
Concept of: single bond – two shared electrons
double bond – four shared electrons
triple bond – six shared electrons

10. 5 Rules for Making Lewis Dot Structures
Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge).
Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required.
Using lone pairs, complete octets around the noncentral atoms.

11. 5 Rules for Making Lewis Dot Structures
4. Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N.
5. If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.

12. Let’s Draw Lewis Dot Structures
Duets and Octets
How would you draw the dot structure for H2 or for CH4?

13. Let’s Draw Lewis Dot Structures
Ammonia, NH3
Required Valence Electrons:
For Atoms
Charge (0)
Total 8
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 6 2

14. Let’s Draw Lewis Dot Structures
Molecular nitrogen, N2
Required Valence Electrons:
For Atoms
Charge (0)
Total 10
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 2 8
   
   
Required Valence Electrons:
For Atoms
Charge (0)
Total 10
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 2 4

15. ? Let’s Draw Lewis Dot Structures Why not Required Valence Electrons:
   
Required Valence Electrons:
For Atoms
Charge (0)
Total 10
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 2 8

16. Let’s Draw Lewis Dot Structures
Carbon dioxide, CO2
Required Valence Electrons:
For Atoms
Charge (0)
Total 16
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 4 12
   
   
Required Valence Electrons:
For Atoms
Charge (0)
Total 16
NONE
Electrons As Written:
Single Bonds
Lone Pair
Double/Triple Bonds
Still Needed 4 8

17. Drawing Lewis Dot Structures
Draw the atoms on paper and put dots around them to represent valence electrons of the atom. Be sure to have the correct number of electrons.
If the species is an ion, add or subtract electrons corresponding to the charge of the ion. Add an electron for every negative (-) charge, and subtract an electrons for every positive (+) charge.
Consider bonding between atoms by sharing electrons; if possible, apply the octet rule to your structure.
NEW: Assign formal charges to atoms in the structure.

18. Formal Charge
Assume each atom satisfies the octet rule (there are exceptions)
Sum of the formal charges for the atoms equals the charge on the molecule (zero) or the molecular ion (minus or plus an integer)
Each atom “owns” ½ of the electrons in each bond and all of the electrons in its non-bonded pairs.
Formal charge = valence electrons – number of electrons “owned”

19. Number of valence electrons corresponds to column number in the periodic table:
H = B = C = N = O,S = F,Cl,Br,I =7
Example: ammonia, NH3 H N .
# owned = 3 x = 5
valence of N = 5
Formal Charge = 5 -5 = 0

20. Another example, NH4+ N H +
Valence of N = 5
Owned by N = 4
Formal Charge on N = = +1
Formal Charge on H = = 0

21. The “correct” Lewis structure has fewer atoms with negative or positive charges. If there is more than one possible Lewis structure, we use calculations of formal charge on each atom to determine the more likely Lewis structure.
For example, take chloric acid, HClO3. Here are two possible structures: or