Unit 5: Chemical
Formulas and Bonding
Chapters 7 and 8
Unit # 5: Chemical Composition
I.
“Intra-Chemical” Bonding
A. Definition: atoms (elements) held
together by an attractive force
B. Types (3)
1. Metallic: atoms of a metal
“share” the valence electrons
because they move from one
element to another
2. Ionic (strongest)
a. Definition: electrons are transferred
from one element to another (opposite
charges is attractive force)
b. Called: “salts”
c. Atom that donates electron:(+)cation
Atom that accepts electron:(-) anion
d. Oxidation State (number):
Charge of an ion
e. Lewis Dot Formulas – used to show outer
shell electron transfer
- Octet Rule: every element wants
8 electrons in its outer shell
1. potassium + chlorine  potassium chloride
(show on board)
2. magnesium + fluorine  magnesium fluoride
(show on board)
f. Types of Ions
1. Monatomic Cation: cation with one
element
2. Monatomic Anion: anion with one
element (ends with –ide)
3. Polyatomic Ions:
Ammonium
Chromate
Acetate
Chlorate
Phosphate
Nitrate
Sulfate
Hydroxide
Carbonate
Permanganate
g. Writing Chemical Formulas
1. Composition
2. Number of elements
3. Writing chemical formulas (from the
name)
a) Recognize the (+) and (-) ions
b) Write the symbols of the
elements with their charge
* Note- a Roman numeral will tell you what the
charge is on the cation if there is more than one
possibility
c) Adjust the number of each ion (with
subscripts) as needed so the positive
charge is equal and opposite the negative
charge.
•
Note – if the ions are polyatomic and there is
more than one, the ion is enclosed with
parentheses with a subscript on the outside.
•
Examples:
Sodium Chloride
Calcium Sulfate
Barium Phosphate
(on board)
(on board)
(on board)
h. Naming Compounds
1. Consists of two words:
a) Name the cation
b) Name the anion
c) If the cation has more than one
possible charge, a Roman Numeral is
used to show the charge.
Ex: FeCl3
FeCl2
+3
Fe
 Iron (III) Chloride
+2
Fe
 Iron (II) Chloride
Ex: NH4Cl
Cu2SO4
NaC2H3O2
3. Covalent (weaker)
a. Definition: valence electrons are shared
between two elements
Ex: F2
(show electron dot diagrams for bonding)
b. Types
1. Polar Covalent (stronger): unequal
sharing of electrons (one pulls more than
the other – the more electronegative
element)
2. Non-Polar Covalent (weakest):
equal sharing of electrons
c. Lewis Structures
1. examples (hydrogen and HCl)
2. The number of covalent bonds formed by
an atom equals the number of unpaired
electrons in the Lewis Dot Formula.
3. Examples
a) water
b) ammonia
c) methane
d. Multiple Bonds
1. Double bonds: two pairs of electrons are
shared (Ex: oxygen gas - draw it)
2. Triple Bonds: three pairs of electrons are shared
(Ex: nitrogen gas- draw it)
e. Hybridization: combining of two or more orbitals of nearly
the same energy into new orbitals of equal energy
Ex. 1) C 1s22s22p2
sp3 hybrid
__ __ __ __ -----> __ __ __ __
2s
2p acquires energy 2s
2p
2) B 1s22s22p1
sp2 hybrid
__ __ __ __ -----> __ __ __ __
2s
2p acquires energy 2s
2p
f. Molecules with more than one element (polar vs. non-polar)
1) Depends On: - electronegativity difference (2 elements)
- Non-bonded electron pairs (2+ elements)
- Structure (symmetry) (2+ elements)
g. Shapes of Molecules (VSPER handout)
h. Naming/Writing Formulas for Covalent Compounds
1. Binary covalent compounds (2 elements)
2. Formulas with two non-metals
3. Rules
a) First word:
- prefix indicating the # of atoms for
the first element (if there is more
than one)
- name of first element
b) Second word:
- prefix for the number of atoms of
the second element
- name of second element
II. “Inter-Chemical” Bonding
A. Definition: whole “salts” or “molecules”
attract and bond with one another.
B. Types
1. Hydrogen Bonding (medium strength):
H bonded to N, O, or F
2. “VanderWaals” Forces
a. Dispersion (weakest): if not H-bonding
or dipole-dipole, it must be dispersion
b. Dipole-Dipole (strongest): between
polar molecules