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Chemical Bonding
• Although we have talked about atoms
and molecules individually, the world
around us is almost entirely made of
compounds and mixtures of compounds.
• We are going to take an in depth look at
these compounds and the interactions
of the atoms that hold them together
and make up the compounds
Bonds• Bonds are the force that holds
groups of two or more atoms
together and makes them function as
a unit.
• Bond Energy is the energy required
to break the bond between two
atoms.
Ionic Bonding
• Generally occurs between a Metal and a
Non-metal
• Cations lose electrons, Anions gain electrons
• Electrons are transferred
• Opposite charges on atoms attracts them to
one another
Covalent Bonding• Generally occurs between a nonmetal and a
nonmetal
• Both atoms share electrons to achieve a
lower energy state.
• Electrons are “shared”. More like a tug of
war.
• Same charges, lower energy is responsible
for bonds
Nonpolar Covalent Bonding
• Covalent bonds where electrons are shared
equally between two atoms.
• Atoms must have the same values of
electronegativity
• If a covalent bond is like a tug of war a
nonpolar covalent bond would be a
stalemate.
Polar Covalent Bonding
• During the “tug of war” in covalent bonding
electrons aren’t always shared equally.
• Some atoms have a stronger attraction for
electrons and pull them closer than other atoms.
• This unequal sharing of electrons causes one
atom to have a small positive charge and one to
have a small negative charge
Electronegativity
• Attraction of shared electrons to an
atom.
• Determines the type of bond
• Can calculate a value of
electronegativity based on relative
values for each element.
Electronegativity• For differences in electronegativity,
generally:
O = nonpolar covalent examples: Cl-
Cl, C-S
.1-1.5 = polar covalent examples: C-F,
P-S
1.6-3.3 = Ionic examples: Na-O, K-I
Practice• Identify the following as Ionic, polar
covalent, or nonpolar covalent bonds:
S-F Mg-Cl
Br-Br B-F
B-N N-Cl
P-I Mn-S
Dipole Moments
• Polar covalent bonds that do not share
electrons equally are said to have a dipole
moment.
• The atom pulling the electrons the strongest
or with the higher electronegativity will have
a partial negative charge.
• The atom with the weaker pull on electrons
will have a partial positive charge.
Lewis Dot Structures
• Representations of atoms or molecules which show
the valence electrons around an atom or molecule
• Hydrogen follows a duet rule – two valence
electrons give it the same electron configuration
as helium
• Most other atoms follow a octet rule – eight
valence electrons will give each atom the same
number of valence electrons as a noble gas
Lewis Dot Structures - Ionic• Metals lose electrons, nonmetals gain electrons
• Rules for LDS for ionic compounds:
1. Write each element symbol
2. Determine the number of valence electrons
3. Add the valence electrons to each atom.
Clockwise- 12,3,6,9 one at a time.
4. Show the electron transfer from metal(s) to
nonmetal(s)
Lewis Dot Structures-Practice• Draw the Lewis Dot Structure for the
following atoms:
• Ca
• F
• Se
• Al
• P
• Si
Lewis Dot Structures - Practice
• Draw the Lewis Dot Structures for the
following ionic compounds:
• Na + Cl
• Mg + Br
• Al + O
• B + F
Lewis Dot Structures -Covalent
• Two nonmetals share electrons to achieve a
lower energy
• Rules:
1-3 same as ionic
4. Circle electrons that will pair together
5. Rearrange the compound so shared electrons
are aligned correctly(between atoms).
Lewis Dot Structures - Practice
• Draw the LDS for the following covalent
compounds:
• CCl4
• NBr3
• Phosphorus + Iodine
• Silicon + Fluorine
Bond Strength
• Of the three covalent bonds:
• A triple bond is the strongest followed by a
double bond and then a single bond which
is the weakest of the three
• A triple bond has the highest bond energy,
then a double bond, followed by a single
bond
Bond Length
• Of the three covalent bonds:
• A triple bond is the shortest followed
by a double bond and then a single
bond which is the longest
• Why?
LDS- Covalent Compounds
• So far we have looked at simple
covalent compounds and how they
will share valence electrons
• Now we will look at more complex
covalent compounds and how to
determine the Lewis Dot Structure.
• Try drawing the Lewis Dot Structure for
SO2 using the rules for covalent
compounds that you have learned
• The actual structure of the compound is:
• Let’s take a look at how to draw LDS
when we are given the formula and the
compound is more complex
Rules for Complex Covalent LDS• 1. Determine the total number of valence
electrons in the compound
• 2. Begin by putting a single bond between each
atom (Choose appropriate middle atom if
necessary)
• 3. Fill in lone pair electrons to fulfill duet/octet
rule
• 4. Add a double bond (or triple bond) if
necessary to insure the duet/octet rules are
fulfilled and the total number of valence
electrons are correct.
Resonance• Resonance occurs when several equally correct Lewis
Dot structures can be assigned to compounds.
• Double arrows are used to show options for
compounds with resonance structures.
• Resonance structures for SO2:
Practice:
• Draw the LDS for the following
compounds with the new rules you
have been given:
HF N2 NH3 CH4
NF3 O2 CO PH3
LDS for Ions
• For ions the rules for drawing Lewis Dot Structures
are the same except the total number of valence
electrons will either increase or decrease
depending on the charge
• For a positive charge subtract a valence electron
• For a negative charge add a valance electron
• After Drawing the LDS brackets are added and the
charge is added outside the brackets- top right
Example: NH4+
• Add up valence electrons for each atom:
N-5 H-1 total = 9
• Because of the +1 charge we assume a
valence electron has been lost
• Our new total of valence electrons is 8
• Draw the LDS using the same rules.
Don’t forget the brackets and the charge
Practice
• Complete the LDS for the following
ions. Show resonance structures if
they exist
NO+ NO3- SO4
-2
ClO3-
PO4-3 SCN-
Lone Pair Electrons• The unshared valence electrons
represented in LDS are called lone
pair electrons.
• Example: In CF4 each fluorine has six
lone pair electrons and
carbon has zero for a total
of twenty four.
VSEPR
• Stands for Valence Shell Electron Pair
Repulsion
• This theory states that electrons pairs
around an atom will spread out as far as
possible
• This repulsion is due to the same charges
on electrons
Molecule Polarity• We have already discussed a polar
covalent bond in terms of dipole
moments caused by differences in
electronegativity.
• We will now use this knowledge to
determine whether a molecule is
polar or non polar
Molecule Polarity
• In order for a molecule to be considered
polar it needs to have a concentrated
partial positive charge on one end and a
concentrated partial negative charge on
the other.
• Molecules that have polar bonds will not
necessarily be polar molecules
Molecule Polarity
• Symmetrical molecules can have
dipole moments cancel each other
out causing them to be nonpolar.
• Examples CH4, BF3, CO2
Molecule Polarity
• A molecule with a lone pair of
electrons in place of a bond will
always be polar (bent and trigonal
pyramidal)
• Examples: NH3, H2O, NO2-
Molecule Polarity
• Molecules that have a tetrahedral,
trigonal planar, and linear geometry
can be either polar or nonpolar.
• The central atom would have to have
different atoms bonded to it to be a
polar molecule. Example: HCN, HCO2-