Chapter 9
Covalent Bonds
Read pgs 240-270
Covalent Bonds
• Covalent bonds form between atoms that
share electrons.
• Covalent bonds form between two or more
nonmetals.
• Shared electrons can be used or counted for
both atoms in a bond.
• A pair of atoms can share more than one pair
of valence electrons.
Multiple Bonds
• If a pair of atoms share 4 electrons, we call it a
double bond.
• Share 6 electrons, a triple bond.
• The more electrons shared, shorter and stronger
bond
• Strength of a bond depends on
– Size of atoms Length of bonds
– Number of shared electrons
Shared bonds between the nucleii of atoms are
called sigma bonds.
Double and Triple bonds form outside of the central
axis and are called Pi bonds.
Naming Covalent Compounds
• The rules for naming covalent compounds are
very similar to ionic.
• Name the 1st element
• Name the 2nd element and change its ending
to –ide.
• For covalent compounds we add a prefix to
show how many of each element that is in the
compound.
• Prefixes
– Mono – 1
– Tri – 3
– Penta – 5
– Hepta – 7
Di or Bi – 2
Tetra – 4
Hexa – 6
Octa – 8
The only exception to the rule is on the first
element.
If you only have one of the first element, you
don’t use the mono- prefix.
Examples
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Name the following covalent compounds
CO
CO2
Carbon Monoxide Carbon Dioxide
NCl3
CCl4
• Nitrogen Trichloride
Carbon Tetrachloride
• SF6
N2S5
• Sulfur Hexafluoride
Dinitrogen Pentasulfide
Naming Acids
• We have already talked about acids.
– Corrosive liquids
– pH’s less than 7.
– Produce H+ ions in water.
• All acids start with hydrogen. You can have
two types of acids: Binary, H and one other
element, or Ternary, H with a polyatomic ion.
• Naming Binary Acids
– Use the prefix Hydro– Name 2nd element and change ending to –ic acid.
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Examples
HCl – Hydrochloric Acid
HF – Hydrofluoric Acid
H3P – Hydrophosphoric Acid
Phosphorus picks up an –or
Sulfur picks up an - ur
• Naming Ternary Acids
– Identify the polyatomic ion.
– If the ion ends in –ate, change ending to –ic Acid.
– If the ion ends in –ite, change ending to –ous Acid.
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Examples
HNO3 – nitrate – Nitric Acid
HNO2 – nitrite – Nitrous Acid
H2SO4 – sulfate – Sulfuric Acid
Lewis Structures
• To show bonding electrons and unshared pairs
of electrons in molecules we can use one of
two systems: Dot Method or Math Method.
• The Lewis structure of a molecule will show
bonds, the structure of the molecule and
eventually the shape of the molecule.
• Choose one method or the other, but don’t try
to do both.
Dot Method
• Find the number of valence electrons of each
element in the molecule.
• Determine the central atom of the molecule.
– The atom with the most single valence electrons or the
element you have the least of.
• Match up single valence electrons to form bonds.
– Coordinate covalent bonds form with a pair of electrons.
• Make sure every element has 8 valence electrons,
(except hydrogen – it only needs two ).
Math Method
• Determine the total number of valence
electrons you have in the molecule.
• Determine how many valence electrons you
need for each element to complete its octet.
• Take the difference between the what you
have and what you need and divide by two to
get the number of bonds that form.
• Choose the central atom and distribute the
bonds and the remaining electrons.
• If an atom needs two electrons and another
atom has a pair of electrons it can form a
coordinate covalent bond where one atom
donates both electrons in the bond.
• Resonance structures are when there is more
than one correct Lewis structure for a
molecule.
• Shapes of molecules play a big part in
determining the chemical and physical
properties of the molecule.
VSEPR Theory
• To determine the shape of a molecule we use
the VSEPR Theory.
• VSEPR stands for Valence Shell Electron Pair
Repulsion
• Bonding electrons have less repulsion than
nonbonding electrons. The nonbonding
electrons push and move the other bonds
around to reduce repulsive forces.
• There can be exceptions to the octet rule.
Molecules can for with atoms having less than
8 valence electrons or more than 8 valence
electrons.
• Because electrons are constantly moving
around the atom it can create momentary
positive or negative areas around the
molecule.
• These weak forces of attraction are called
dispersion or Van der Waal’s forces.
Polar Bonds and Molecules
• Stronger attractions are created when the
positive/negative poles in a molecule are
more constant.
• These dipoles are created when there is
unequal sharing of electrons in a covalent
bond.
• To determine if a bond is polar look at the
difference in electronegativities.
• If the difference in E.N. is less than .4, than the
bond is nonpolar. The electrons are shared
equally, no force of attraction.
• If the difference in E.N. is between .4 and 2.0,
then the bond is polar. The molecule is
dipolar, one end is always negative.
• If the difference is greater than 2.0, then the
bond is ionic, electrons are transferred.
• Diff < 0.4
• 0.4 ≤ Diff. ≤ 2.0
• Diff. > 2.0
nonpolar – equal sharing
polar – unequal sharing
ionic – transfers electrons