Chemical Bonds
Forming Chemical Bonds
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The force that holds two atoms together is called a
chemical bond .
The valence electrons are the electrons involved in
forming chemical bonds.
Elements tend to react to acquire eight electrons.
This is call a stable octet.
Noble gases (group VIIIA/18) have this structure
(octet) and are inert (does not form bonds).
Atoms can gain, lose, or share electrons to reach
an octet.
Forming Ions
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Positive ions (cations) are formed when atoms
lose one or more valence electrons. [Usually
these atoms are metals.]
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Reactivity of metals (cations) are based on how
easily they lose electrons. [ Ionization energy]
Negative ions (anions) are formed when atoms
gain one or more valence electrons. [These
atoms tend to be non-metals.]
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Electronegativity is the ability to attract or gain
electrons.
Forming Ionic Bonds
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Ionic bonds: Complete transfer of electrons
between atoms (difference of electronegativity of
1.6 or greater).
Two neutral atoms will form ions. The resulting
compound is called an ionic compound.
Properties of Ionic Compounds
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Form crystal lattice bonding structure.
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Example: NaCl (sodium chloride)
High melting point and boiling point due to
strong electrostatic charge between the atoms
(cation is attracted to anion after the transfer of
electrons).
Hard, rigid and brittle solids.
Conducts electricity in liquid state or
dissolved in water only.
Forming Covalent Bonds
Covalent bonds: Sharing of electrons between
atoms (difference of electronegativity of less
than 1.6).
 Occurs usually between elements close to each
other on the periodic table (mostly nonmetals).
 The resulting compound is called a molecule.
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Properties of Covalent Compounds
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Have definite and predictable shapes.
Low melting and boiling points.
Relatively soft solids.
Can exist as solids, liquids or gases.
Comparison of Bonding Types
ionic
covalent
ions
valence
molten salts
electrons
conductive
transfer of electrons
high mp
DEN > 1.6
molecules
sharing of electrons
non-conductive
low mp
DEN < 1.6
Determining Number of
Covalent Bonds
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To determine how many bonds exist in a
molecule, use the following formula
N - A = # bonds
2
-- Where N is the # of needed electrons, which is 8
for all elements but H, which is 2.
-- Where A is the # of available electrons, which is
the number of valence electrons.
Lewis Dot Structures
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A system of “drawing”
bonds
Shows how valence
electrons are arranged
Dots represent valence
electrons
Pairs of dots represent
bonding pairs of
electrons.
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Show the reaction between
sodium and chlorine:
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Show the reaction between
calcium and bromine:
Drawing Lewis Structures
1.Determine the # of bonds need with the formula
2.Draw a skeleton structure. The least electronegative
element is the central atom.
3.Connect the atoms with the # of bonds determined in
step 1.
4.Finish by making sure all atoms in the structure have
an octet of electrons.
5.Remember, if the structure has a positive charge, it
has fewer available electrons, and if it has a negative
charge it has more available electrons.
Lewis Structure Practice
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Draw Lewis Dot Structures for the following
molecules and polyatomics:
CO2 N2 O2
Cl2
H2O SO4 -2
Exceptions to the Octet Rule
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There are two elements which don’t want an octet of electrons
and those are Be (wants 4) and B (wants 6).
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Third-row and heavier elements often satisfy the octet rule but
can exceed the octet rule using their empty valence d orbitals,
especially when surrounded by highly electronegative atoms
such as chlorine, bromine, and oxygen.
Exceptions to the Octet Rule
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The exceptions are easy to recognize-- when you
calculate the # of bonds, the formula will give you an
answer that makes no sense. If this happens, do the
following:
1.
2.
3.
Connect all of the atoms to the central element with one
bond.
Give all of the surrounding elements an octet of electrons.
Count the number of electrons you have put in the structure
so far.
- If it equals the available # of electrons you calculated, you’re done.
- If it is less than than the available # of electrons you calculates, place
the needed number of electrons around the central element.
Exceptions to the Octet Rule
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Draw Lewis Dot Structures for the following
molecules:
BF3
PCl5
SF6
XeF2
Carbon Bonding
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Carbon follows the rules previously discussed, but it
has a few unique properties.
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Carbon is always the central atom in its compounds
and it tends to undergo bonding to itself. This
tendency of carbon to bond to itself is called
catenation.
Carbon always has four bonds.
Anytime you have a choice for a skeleton structure,
always go for the most symmetrical option.
Carbon Bonding
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Draw Lewis Dot Structures for the following
molecules:
C2H6
C3H8
C2H2
Let’s Practice!
H2
 SO2
 CCl4
 SiH4
 PCl5
 XeCl6
 H 2O
 SO4-2
 NO3 C2H3O2
Molecular Shape
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Valence Share Electron-Pair Repulsion
(VSEPR) model allows us to predict the
molecular shape by assuming that the
repulsive forces of electron pairs cause them
to be as far apart as possible from each other.
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Only the valence electron pairs are considered in
determining the geometry.
Effect of the number of electron pairs
around the central atom
2 charge clouds,
linear
3 charge clouds,
trigonal planar
5 charge clouds,
trigonal bipyramidal
4 charge clouds,
tetrahedral
6 charge clouds,
Octahedral
PREDICTING EXPECTED GEOMETRY
ACCORDING TO VSEPR THEORY
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Lewis dot structure determines the total # of electrons
around the central atom. Multiple bonds (double and
triple) count as one.
The number of bonding and nonbonding electron pairs
around the central atom determines the geometry of
electron pairs and the molecular geometry.
Lone e pairs affect geometry more than bonding pairs.
Multiple bonds have larger affect on geometry than single
bonds: H2C=O (116° instead of 120°); H2C=CH2 (117°
instead of 120°).
Molecular Shapes 2,3,4 Electron Pairs
Molecular Shapes 5, 6 Electron Pairs