3. Definition
• Valence Electrons: are electrons in the
outermost shell (energy level).
– The Electrons in the S-Block and P-Block
– DOES NOT INCLUDE electrons in the D-Block
and F-Block
• They are only filled after a new valence shell has been
started.
• They are the electrons available for
bonding.
• The number of valence determines what
type of bonds will usually be formed
14. Lewis Dot Diagrams
Lewis Dot Diagrams are a tool to help you visually represent the
valence shell.
Note: Lewis Dot structures follow Hund’s Rule and the Pauli Exclusion Principle
1 valence e- 2 valence e- 3 valence e- 4 valence e-
X X X X
5 valence e- 6 valence e- 7 valence e- 8 valence e-
X X X X
15. Dot Notations: Period 2
Lewis dot notations for the valence electrons of
the elements of Period 2.
lithium beryllium boron carbon
Li Be B C
nitrogen oxygen fluorine neon
N O F Ne
16. The Octet Rule
• Octet Rule: Atoms tend to gain, lose, or share
electrons in order to gain a full set of 8 valence
electrons.
– Want to be like Noble Gas
– What about Period 1?
• When atoms gain/loose electrons they become
ions.
– Anion: Gains Electrons (Net – Charge)
– Cation: Looses Electron (Net + Charge)
• Oxidation Number: Represents the charge of a
typical ion of the element.
– Tells you how many electrons are typically gained or
lost.
17. Bonding
• Chemical Bond: Any force that holds two
atoms together.
– Created when two (or more) atoms interact with
one another in a way that fulfills the Octet Rule.
• Two or more different Elements bond
together to form Compounds.
– Compounds can have very different chemical and
physical properties than their component
elements.
24. Ionic Bonding
• Ionic Bonding: Occurs when one atom donates an
electron and another receives the electron.
• The exchange (gain/loss) of an electron creates an
Anion (- Charge) and Cation (+ Charge).
– Remember opposites attract!
• The electrostatic attraction between the Anion and Cation
causes them to stick together forming an ionic bond.
– Also known as electrostatic bonds.
• 1 or more electrons are actually exchanged
between the two atoms.
• Occur when there are large differences in ionization
energy and electronegativity.
– Between metals and non-metals.
25. Examples of Ionic compounds
Mg2+Cl-2 Magnesium chloride: Magnesium loses two
electrons and each chlorine gains one
electron
Na+2O2- Sodium oxide: Each sodium loses one
electron and the oxygen gains two electrons
Al3+2S2-3 Aluminum sulfide: Each aluminum loses
two electrons (six total) and each sulfur
gains two electrons (six total)
26. Metal Monatomic Ion name
Cations
Lithium Li+ Lithium
Sodium Na+ Sodium
Potassium K+ Potassium
Magnesium Mg2+ Magnesium
Calcium Ca2+ Calcium
Barium Ba2+ Barium
Aluminum Al3+ Aluminum
28. Sodium Chloride Crystal Lattice
Ionic compounds form solid
crystals at ordinary
temperatures.
Ionic compounds organize
in a characteristic crystal
lattice of alternating
positive and negative ions.
These ionic bonded crystals are known as salts.
29. Properties of Ionic Compounds
Structure: Crystalline solids
Melting point: Generally high (strong
bonds)
Boiling Point: Generally high (strong
bonds)
Electrical No conductivity in solid
Conductivity: form. Excellent conductors
when molten and aqueous
(dissolved in liquid).
Solubility: Water: Soluble
Alcohol: Insoluble
31. Covalent Bonding
• What happens when both atoms want to gain electrons?
– Covalent Bonding: Bond formed when two atoms share one or more
valence electrons to achieve full octets.
– Occurs when both have have high ionization energies and
electronegativity.
• Both try to take electron from each other… tug-of-war.
• Non-Metal to Non-Metal
• Typically less strong than ionic bonds.
• The term “molecule” is used exclusively for covalent bonding
F F
32. The Octet Rule:
The Diatomic Fluorine Molecule
F Each has seven
1s 2s 2p
valence electrons
Each Wants 1 More
F Electron
1s 2s 2p
F F
Single Bond: 1 Electron is shared by each atom (total of 2 shared electrons).
33. The Octet Rule:
The Diatomic Oxygen Molecule
O Each has Six
1s 2s 2p valence electrons
Each Wants 2 More
O Electrons
1s 2s 2p
O O
Double Bond: 2 Electrons are shared by each atom (total of 4 shared electrons).
34. The Octet Rule:
The Diatomic Nitrogen Molecule
N Each has Five
1s 2s 2p
valence electrons
Each Wants 3 More
N Electrons
1s 2s 2p
N N
Triple Bond: 3 Electrons are shared by each atom (total of 6 shared electrons).
35. Representing Covalent Bonds
Covalent Bonds (shared electrons pairs) can be
represented by two dots (:) or by a single line ( - )
O= O
O O
36. Properties of Covalent Molecules
Structure: Molecular Chains
Melting point: Generally low (weak bonds)
Boiling Point: Generally low (weak bonds)
Electrical Poor conductors (no free
Conductivity: ions or electrons)
Solubility: Water: Insoluble
Alcohol: Soluble
38. Metallic Bonding
• Metals tend to be electron donors.
• When metals bond with other metals
their valence electrons become
delocalized (free to move around)
and form what is known as an
electron sea.
o Valence electrons do not belong to
any one atom.
• A metallic bond results from the
attraction between the metal cations
and the surrounding sea of
delocalized electrons
39. Packing in Metals
Lattice Model: Packing
uniform, hard spheres to
best use available space.
This is called closest
packing. Each atom has
12 nearest neighbors.
• Vacant p and d orbitals
in metal's outer energy
levels overlap, and
allow outer electrons
to move freely
throughout the metal
40. Metal Alloys
Substitutional Alloy:
some metal atoms
replaced by others of
similar size.
Interstitial Alloy:
Interstices (holes) in
closest packed metal
structure are occupied
by small atoms.
41. Properties of Metals
Metals are good conductors
of heat and electricity (lots of
free electrons).
Metals are malleable
Metals are ductile
Metals have high tensile
strength
Metals have luster
