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
NF3
N2
NH3
O2
CH4
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+
PO4-3
NO3SCN-
SO4-2
ClO3-
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-