Chemical Bonds
We love them…we really
do!
Exactly what are
chemical bonds???
» Defined as: a mutual electrical attraction
between the nuclei and valence
electrons of different atoms that binds
the atoms together
» The reason bonds are formed is
because they are more stable bonded
together than as individual atoms.
Types of Chemical
Bonds
»1) Ionic bonds
»2) Covalent bonds
i. polar covalent
ii. non-polar covalent
»3) Metallic bonds
Ionic Bonds
Characteristics:
1) Results when electrons are completely
given from one atom to another
2) Occurs between metals and nonmetals
Covalent
Bonds
» Characteristics:
» 1) Results from the sharing of electron
pairs between two atoms
» 2) Occurs when nonmetals bond to
each other
Ionic or Covalent?
» The most important thing to
remember about bonds is
that they are on a
continuum from ionic to
covalent.
» This means that you can
have more ionic or more
covalent character. That is
why we use % ionic
character.
Electronegativity
Differences
» To use the EN chart » If it falls the in the
on page 162 of your
range of 0-0.3 then
text:
it is non-polar
covalent.
» Find the difference in
EN of the two atoms » If it falls in the range
involved in the bond.
of 0.4-1.6, it is polar
(you can find EN
covalent.
values on the new
» If it falls in the range
periodic table I gave
of 1.7-3.3, it is ionic.
you!)
Non-polar Covalent
Bonds
» These are a type of covalent bond in
which the electrons are shared equally
between the two atoms in the bond.
» This means that the negative charge is
distributed evenly!
» Examples of compounds that are nonpolar: hexane and oil
Polar Covalent Bonds
» Polar covalent bonds do not share
electrons equally between the two
atoms involved.
» This means that “poles” of positive and
negative charges begin to form at each
atom.
» Examples of polar compounds: water,
ethanol
Molecular compounds
» A molecule is a neutral group of atoms
that are held together by covalent
bonds.
» Examples: H2O and CO2
» A chemical compound whose simplest
units are molecules is called a
molecular compound.
Chemical vs. Molecular
Formulas
» Really, chemical formulas
encompasses all the formulas involving
chemicals (for all types of bonding).
» Molecular formulas are types of
chemical formulas that describe
molecules held together by covalent
bonds.
Diatomic Molecules
» Molecules that contain only two atoms
are diatomic.
» There are several molecules that exist
as diatomics in nature.
» HOFBrINCl - (pronounced hoffbrinkle)
Formation of a
covalent bond
» When atoms at a distance are first attracted
to one another, the attraction of their nucleus
with the other atoms electrons is STRONGER
than the repulsion between the nuclei and
electrons.
» Therefore, they will come close enough
together to bond. When they reach the
optimal distance, their potential energy is at
its lowest.
Characteristics of a
Covalent Bond
» The distance between two bonded atoms at
their minimum potential energy is called the
bond length.
» Thus each atom releases energy as they
change from individual atoms to a molecule.
» The same amount of energy must be
supplied to separate the newly formed bond.
This energy is called the bond energy.
Bond energies
» Scientists report these energies in
kJ/mol.
» To break a H-H bond, 436 kJ of energy
is needed.
» Bond energies and bond lengths are
different depending on which two
atoms are involved in the bond.
The Octet Rule
» Chemical compounds tend to form so
that each atom, by gaining or losing
electrons, has an octet of electrons in its
highest occupied energy level.
Exceptions to the Rule
» Hydrogen- can only bond to one other
atom…wants 2 valence electrons.
» Boron wants 6 valence electrons. (It already
has 3!)
» Some elements can be surrounded by MORE
than 8 electrons (called expanded valence)
when bonded to halogens or other highly EN
elements. Examples: P, As, S
Electron Dot Notation
» Shows the valence electrons of an element
» Uses the element symbol surrounded by up
to 8 dots.
» The order in which to place the dots:
» Generally, one dot is placed on each side first
before pairing the dots.
» Let’s do some examples!
Lewis Structures
»Electron dot notation can also
be used to represent molecules.
» In this case, the valence
electrons also show us how the
atoms are bonded.
Lewis Structures
» Unshared pairs of
electrons are called
lone pairs. These are
electrons that are not
involved in bonding.
» A pair of dots involved
in bonding may also be
represented as a dash.
Structural
Formulas
» A structural formula
indicates the kind,
number,
arrangement and
bonds but does not
include the unshared
or lone pairs of
electrons.
How to Write a Lewis
Structure
» 1. Determine the
type and number of
atoms in the
molecule.
» 2. Add together the
total number of
valence electrons
involved.
» 3. Arrange the
atoms to form a
skeletal structure of
the molecule. If
carbon is present, it
is always in the
center!
Continued
» 4. Add unshared
pairs of electrons
where appropriate.
» 5. Count the
number of electrons
to be sure that the
number used equals
the number
available.
»
»
»
»
»
Practice:
CH3I
NH3
H2S
CO3-2
Multiple covalent bonds
» Some elements can
share more than
one pair of
electrons, especially
carbon, nitrogen,
and oxygen.
» A double bond is a
covalent bond
produced by sharing
two pairs of
electrons.
» A triple bond is a
covalent bond
produced by sharing
3 pairs of electrons!
About Multiple Bonds
» Double and triple
bonds have higher
bond energies and
have shorter bond
lengths than single
bonds.
»
»
»
»
Practice:
N2
HCN
CO2
Resonance Structures
» Some molecules
and ions cannot
accurately be
represented with
one Lewis structure.
» This occurs when a
molecule is
asymmetrical with
respect to bonds of
the same type.
» Example:
» Ozone
» When writing
resonance
structures you must
include all the
possibilities.
Covalent Network
Bonding
» All the covalent
molecules you have
learned about to this
point are molecular.
» Some covalent
molecules do not exist
as individual molecules.
They are bound
together by forces
acting between them.
» Continuous 3-D
networks of bonded
atoms are referred
to as a covalent
network. You will
learn more about
these later!
More About Ionic
Compounds
» Ionic compounds are composed of
positive and negative ions that are
combined so that the number of
positive and negative charges are
equal.
» Most ionic compounds exist as
crystalline solids. Many minerals are
ionic compounds.
Formula Units
» The chemical formulas of ionic
compounds show the ratio of ions
present in any size sample.
» The ratio of ions depends on the
charges of each.
» Example: Ca and F
Characteristics of Ionic
Bonding
» Ions in ionic compounds often form a
crystal structure of repeating units.
» The 3D arrangement of ions depends
on the strength of attraction between
them and their sizes.
» To compare bond strengths, chemists
use lattice energy.
Lattice Energy
» The energy released from one mole of
an ionic crystalline compound as it turns
to a gas is called the lattice energy.
» The values are negative indicating that
the energy is being released from the
compound.
Intramolecular vs.
Intermolecular Forces
» Intramolecular forces are forces IN a
molecule that keep them together.
» Intermolecular forces are those
BETWEEN molecules.
» The difference between these two forces
is why ionic and molecular compounds
have such different physical properties.
Physical properties
» Predict the physical properties
regarding the following substances with
respect to melting point, solubility and
electrical conductivity, in and out of
water:
» Sugar:
» Salt:
» Copper metal:
Melting point
» For ionic compounds, melting points are high.
(Salt)
» This is because the ions are so strongly
attracted to each other in a closely compact
crystalline solid.
» For molecular compounds, the force between
molecules is not as high. This is why they
melt so much easier. (Sugar)
Electrical Conductivity
» In the solid state, the ions cannot move
in an ionic compound, so they are not
good conductors of electricity. (NaCl)
» When they are dissolved in water, they
conduct electricity very well because the
charges are able to move around. (salt
water solution)
More Electrical
Conductivity
» If the solid is molecular (covalent
bonds), it does not develop strong ions
as it dissolves so it would not be a good
conductor. (sugar)
» If the solution has a molecule dissolved
in it, it would not be a good conductor of
electricity because the molecule will not
produce ions with an electrical charge.
(sugar)
Polyatomic ions
» This group of molecules have atoms
that bond together covalently with
each other to form a group that
acts like an ion.
» Examples: OH- ClO3- PO4-3
Metallic Bonding
» Bonding is WAY different in metals than in
ionic and molecular compounds.
» The vacant orbitals in the valence energy
levels overlap with other metal atoms.
» This allows electrons to roam freely between
the atoms. (copper)
» The electrons are said to be delocalized.
» The attraction that results from the attraction
of metals and their sea of electrons is called
metallic bonding.
Characteristics of Metal
Bonds
» The freedom of electrons is what makes
metals a good conductor of electricity!
» The reason that metals are shiny and
reflect light is because the metals can absorb
a wide range of frequencies. The valence
electrons get excited and move to a higher
energy level. When they return to their
ground state light is emitted!
The Strength of Metallic
Bonds
» Metallic bond strength is measured by
the amount of heat required to vaporize
the metal (called the heat of
vaporization).
» It depends on the nuclear charge and
the number of electrons!
Molecular Geometry
» This is the 3D arrangement of
molecules in space.
» The polarity and geometry of the
molecule determine molecular polarity,
or the uneven distribution of molecular
charge.
2 Theories about
Geometry
» 1) VSEPR- (pronounced vess-per)
» Stands for Valence Shell Electron Pair
Repulsion
» The theory is based on the idea that
repulsion between sets of valence level
electrons surrounding atoms causes
them to be oriented as far apart as
possible.
Types of Geometry
» Linear- occurs when two atoms are
bonded to a central atom
» The atoms are 180° apart.
» Example:
Types of Geometry
» If there are 3 identical atoms
surrounding a central atom, then to
orient them as far apart as possible, the
bond angle would be 120°.
» The shape would be trigonal-planar.
Types of Geometry
» Molecules that have 4
atoms bonded to the
central atom always
follow the octet rule.
» The geometry here
would require the angles
to be 109.5°.
» The shape would be
tetrahedral as a result.
Do lone pairs have any
effect of geometry?
» Short answer…YES!
» VSEPR theory says that the lone pair
occupies space around the molecule
just like bonded atoms do.
» This means that there are more shape
possibilities for molecules with lone
pairs!
More Geometry
» For example,
» While the lone pairs
ammonia is NH3 and
do take up space, we
has 3 bonds and 1
describe the shape of
the molecule with
lone pair. It
respect to the position
geometry is
of the atoms ONLY!
described as
trigonal pyramidal.
Bent
» A molecule that
has 2 bonds
and 2 lone
pairs on the
central atom is
considered
bent.
» The bond angle
is about 105°.
» Water is an
example of a
molecule that is
bent.
Trigonal Bipyramidal
» This type of
molecule results
from 5 bonded
atoms with no lone
pairs.
» There are bond
angles of both 90°
and 120°.
» Example: PCl5
Octahedral
» Example: SF6
» In this molecule 6
atoms are bonded to
a central atom with
no lone pairs.
Hybridization
» 2)Hybridization of
molecular molecules is
the 2nd way to predict
shape.
» Hybridization is the
mixing of two or more
atomic orbitals of
similar energies to
produce new orbitals
of equal energy.
» Types of hybrid
orbitals:
» sp
» sp2
» sp3
Geometry of Hybrid
Orbitals
Atomic
orbitals
Type of
Hybrid
Geometry
sp
# of
Hybrid
Orbitals
2
s,p
s,p,p
sp2
3
Trigonal
planar
s,p,p,p
sp3
4
tetrahed
ral
linear
Intermolecular Forces
» Remember…these are forces between
covalent molecules!
» They vary in strength but are weaker
than covalent, ionic and metallic bonds.
» Boiling point is a good measure of
intermolecular forces!
Molecular Polarity
» Dipoles are created by equal but
opposite charges that are separated by
a short distance.
» The direction of a dipole is from positive
to negative pole.
» The forces of attraction between polar
molecules are called dipole-dipole
forces.
Molecular Polarity
» Dipoles determine the polarity of a
bond, and in molecules with more than
one bond, ALL dipoles and their
directions must be considered in
determining molecular polarity.
Examples
» CCl4
» Carbon Dioxide and
formaldehyde
Dipole-Dipole
Interactions
» Short range, only affect molecules that are
near one another.
» Results from the polarity of bonds in a
molecule.
» The partially positive atom attracts to the
partially negative atom in another molecule.
» This force of attraction keeps them together.
Hydrogen Bonding
» Example: Water
» A particularly strong
dipole-dipole force in
which the hydrogen
attached to a highly
electronegative atom is
attracted to the
electronegative atom
in a nearby molecule.
London Dispersion
Forces
» London dispersion
forces are
intermolecular
attractions resulting
from the constant
motion of electrons.
» They are present
between all atoms
and molecules.
» London forces are
the only
intermolecular
forces involved
between noble
gases and non-polar
molecules.
» Strength increases
with increasing
atomic mass of
atoms involved.
Important Things to
Remember
» The important thing to
remember about
intermolcular forces is
that the strength of the
forces between
molecules can help us
predict physical
properties like boiling
point and surface
tension.
» Also, the difference
in electronegativity
of atoms involved in
the intermolecular
force matters
immensely.