Chemical Bonding
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Bonds form in 2 main ways
atoms share electrons
electrons are transferred between
atoms
Type of bond depends on the
atom’s electronegativity and
electron configuration
3 Main Types of Bonds
1. Ionic Bonds

electrostatic force

atoms transfer e- to become ions

usually between a metal and a nonmetal
2. Covalent Bonds

electrons are shared by atoms

usually between two nonmetals
3. Metallic Bonds

forces that hold metals together

metals have many freely moving
electrons that attract positive metal
ions

Ex. What type of bonding would
exist in solid aluminum?
Ionic Bonding
valence electrons: outermost s and p
electrons of an atom
Dot Diagrams: show valence electrons
Examples:
isoelectronic: having the same electron
configuration
Characteristics of Ionic
Compounds
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high melting point
able to conduct electricity in molten
state
tend to be water soluble
crystallize in definite patterns (crystal
lattice)
A closer look at ionic bonding
Naming BINARY Ionic
Compounds

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name the metal first – do not change
ending
name nonmetal second – change
ending to –ide
Examples:
Writing Formulas for Ionic
Compounds
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Sum of all ion charges MUST equal
ZERO!
Use the “criss-cross” method
Examples:
Covalent Bonding
molecule: name for a covalently bonded
particle
Characteristics of Covalent
Compounds

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low melting point
do not conduct electricity
usually brittle solids, liquids, or gases
A closer look at covalent bonding
Naming Covalent Molecules

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make sure the bond is covalent (usually 2
nonmetals)
first element’s name does not change
second element’s ending becomes –ide
Use Greek prefixes to indicate the # of atoms
of each element
mono = 1
penta = 5
di = 2
hexa = 6
tri = 3
hepta = 7
tetra = 4
octa = 8
Writing Formulas for Covalent
Molecules

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Prefixes tell the # of atoms of each
element
Examples:
Molecular Geometry
VSEPR Theory
Lewis Structures


Show arrangement of atoms in
molecules
Show shared (bonding) and free
electrons
Drawing Lewis Structures
Used for covalently bonded molecules
ONLY!
1. Determine the atoms in the molecule
2. Count valence electrons for each atom.
3. Find total # of valence electrons
4. Arrange atoms in skeleton structure.
*Least electronegative atom in center!*
5. Add electrons to structure.

The number of covalent bonds normally
formed by an atom in a Lewis structure
depends on its group in the periodic table.
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H is expected to form one bond.
F, Cl, Br, I, all in group 17 are expected to form
one bond each.
O, S, Se, in group 16, are expected to form two
bonds each.
N, P, As, in group 15, are expected to form three
bonds each.
C, Si, Ge, in group 14 are expected to form four
bonds each.
Octet Rule
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Atoms try to achieve Noble Gas
configuration (8 outer e-)
Hydrogen – forms “duet” instead
Some atoms exceed octet – more than
8 bonding e-
VSEPR Theory
Valence Shell Electron Pair
Repulsion Theory

electron groups arranged to minimize
repulsion
Molecular Shapes

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FILL IN SHAPE CHART!
Show relative positions of atomic nuclei
MUST determine Lewis structure to
determine shape!
Resonance

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Equivalent Lewis structures
Shows possible locations of double
bonds
Resonance
Polarity of Bonds
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polar: having opposite ends
polar bond: e- shared unequally
 caused by difference in
electronegativity
nonpolar bond: e- shared equally
ALL COVALENT BONDS ARE
POLAR EXCEPT:
1.
2.
C–H
any atom bonded to itself
A closer look at polar bonds
Polarity of Molecules
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Bonds must be polar for molecule to be
polar.
Molecule must have a definite top and
bottom with opposite charges in order to
be polar.
Intermolecular Forces (Weak Bonds)
Three main types
1.
Dispersion forces (London, van der
Waals)
2.
Dipole-Dipole Interactions
3.
Hydrogen bonding
Dispersion Forces (van der Waals)
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Very weak
Between nonpolar molecules
Induces momentary (temporary) dipole
Ex. – occurs in Cl2, CO2, CH4, etc.
Dipole-Dipole Interactions
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Stronger than dispersion
Occur between molecules
with permanent dipoles
(aka – polar)
Partially + end of one
molecule attracted to
partially – end of another
Hydrogen bonding

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
Stronger type of dipole-dipole
interactions
Results from H being covalently
bonded to either F, O, or N
Stronger because…
 H is so small
 F, O, & N are very EN
 Partial +/- charges are more
intense
Intermolecular Forces
H-bonds > Dipole-dipole > Dispersion
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Affect BP, MP, solubility
More E required to boil/melt
substances w/ stronger
intermolecular forces
 Why?
Pop Quiz
1. Name CaCO3
2. Write a formula for sodium sulfite.
3. Draw a dot diagram for boron (B).
4. How many valence electrons does
carbon have?
5. What is the oxidation number of
potassium?