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									Chapter 6 -Chemical Bonding
                 A steroid alkaloid derived
                 from skin secretions of
                 the Phyllobates and
                 Dendrobates genera of
                 South American poison-
                 arrow frogs. It is one of
                 the most potent venoms
 Forces that hold groups of atoms together and
 make them function as a unit.

  Ionic bonds– transfer of electrons

 Covalent bonds– sharing of electrons

 Polar Covalent bonds – unequal sharing of electrons
                      that results in an unbalanced
                      distribution of charge
                   The ability of an atom in a molecule
                    to attract shared electrons to

Linus Pauling
 1901 - 1994
Table of Electronegativities
           Shielding Effect
• Electrons in the inner energy levels block
  the attraction of the nucleus for the
  valence electrons.

• Shielding increases down a group.

• This causes EN values to decrease.
          What is Nuclear Shielding?

                     e-                The nucleus (+)
                                       pulls the
                                       electron (-)
                                       close to the
Inner electrons                        core.
shield outer                      e-
electrons from
                        P+ P+
the nuclear pull.    P+ +    P+           The further the
                        P  P+
                              +           electron is away
                       P+ P+ P            from the nucleus,
                                          the weaker the
                                          nuclear pull.

  Range of EN Values                                 ∆EN Values

3.3                  2.0                    0.5                 0

      Mostly Ionic         Polar Covalent         Mostly Covalent

                     Polar covalent- Electrons
                     are shared, but unequally.
                      There is some degree of
                      ionic character in these
Nonpolar          Polar
Covalent         Covalent         Ionic

                                          EN ∆
           0.3              1.7
           Practice Problems
    Determine the bond type using EN values:
         NaCl                      H2O

Sodium’s EN = 0.9        Hydrogen’s EN = 2.1
Chlorine’s EN = 3.0      0xygen’s EN = 3.5

3.0 - 0.9 = 2.1          3.5 - 2.1 = 1.4
therefore it is mostly   therefore it is polar
  ionic                    covalent
  Chapter 6

6.2 Covalent Bonds
Covalent Bonds
             Covalent Terms
• Molecule: A neutral group of atoms that are held
  together by covalent bonds
• Diatomic Molecule: A molecule containing only two
                Example: BrINClHOF
• Molecular Compound: A chemical compound whose
  simplest units are molecules
• Chemical Formula: Indicates the relative numbers of
  atoms of each kind of a chemical compound by using
  atomic symbols and numerical subscripts
• Molecular Formula: Shows the types and numbers of
  atoms combined in a single molecule of a molecular
Bonding Forces

         Electron – Electron
         repulsive forces
         Proton – Proton
         repulsive forces
         Electron – Proton
         attractive forces
Pure Covalent Bonding
Bond Length Diagram
            Bond Energy

It is the energy required to break a bond.
It gives us information about the strength of a
bonding interaction.
Electron Dot
               The Octet Rule
Chemical compounds tend to form so that each atom,
by gaining, losing, or sharing electrons, has an octet
of electrons in its highest occupied energy level.

                   Diatomic Fluorine
Hydrogen Chloride by the Octet
Formation of Water by the Octet
Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet
2nd row elements B and Be often have fewer than
8 electrons around themselves - they are very
3rd row and heavier elements CAN exceed the
octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first,
then place electrons around elements having
available d orbitals.

Shows how valence electrons are arranged
among atoms in a molecule.
Reflects central idea that stability of a compound
relates to noble gas electron configuration.
    Completing a Lewis Structure
  Make carbon the central atom
 Add up available valence electrons:

      C = 4, H = (3)(1), Cl = 7 Total = 14

  Join peripheral atoms to the central
atom with electron pairs.                  .. ..
                                         H C Cl

                                           ..     ..
  Complete octets on atoms other              H
than hydrogen with remaining
    How to Make a Lewis Dot
    Structure for a Molecule
• Draw the Lewis Dot Diagram for each element
• Place the atom with the lowest EN value in the
• Attach the rest of the atoms to the central
• If lone electrons exist on adjacent atoms, pair
  them for multiple bonds.
        Multiple Covalent Bonds:
             Double Bonds

•Two pairs of shared electrons

•Higher bond energy and shorter bond length than
single bonds
       Multiple Covalent Bonds:
             Triple bonds

•Three pairs of shared electrons
•Higher bond energy and shorter bond length than
single or double bonds
 Occurs when more than one valid Lewis
 structure can be written for a particular

 These are resonance structures.
 The actual structure is an average of
 the resonance structures.
Resonance in Ozone

Neither structure is correct.
 Resonance in Polyatomic Ions
Resonance in a carbonate ion:

Resonance in an acetate ion:
   Covalent Network Compounds
Some covalently bonded substances DO NOT form
              discrete molecules.

   Diamond, a network     Graphite, a network of
   of covalently bonded   covalently bonded
   carbon atoms           carbon atoms
  Models are attempts to explain how nature operates
  on the microscopic level based on experiences in
  the macroscopic world.
Models can be physical as
with this DNA model

Models can be mathematical

Models can be theoretical or
Fundamental Properties of Models
 A model does not equal reality.
 Models are oversimplifications, and are therefore
 often wrong.
 Models become more complicated as they age.
 We must understand the underlying assumptions
 in a model so that we don’t misuse them.
 Chapter 6

6.3 Ionic Bonding
Ionic Bonds
Ionic Bonds
    Electrons are transferred

 Electronegativity differences are generally greater
than 1.7
 The formation of ionic bonds is always exothermic!
               Ionic Bonding
• Formula Unit: The simplest collection of atoms from
  which an ionic compound's formula can be
• Lattice Energy: The energy released when one mole
  of an ionic crystalline compound is formed from
  gaseous ions
          Na+ (g) + Cl- (g) → NaCl (s) + 787.5 kJ

• Formation of ionic compounds is ALWAYS
      Sodium Chloride Crystal Lattice
Ionic compounds form solids
at ordinary temperatures.

Ionic compounds organize in
a characteristic crystal lattice
of alternating positive and
negative ions.

Formation of sodium chloride
Na = 3s1 Cl = 3s23p5
Na+ = 2s22p6 Cl- = 3s23p6
            Polyatomic Ions
• A charged group of covalently bonded atoms
• Creation of octets results in an excess or deficit
  of electrons
A Comparison of Ionic and
  Molecular Compounds
  Chapter 6

6.4 Metallic Bonding
      The Metallic Bond Model
• The chemical bonding that results from the
  attraction between metal atoms and the
  surrounding sea of electrons
• Electron Delocalization in Metals
• Vacant p and d orbitals in metal's outer energy
  levels overlap, and allow outer electrons to move
  freely throughout the metal
• The valence electrons do not belong to any one
     Swim in the Sea of Valence
• In metals, the valence electrons are held
  loosely. The vacant p & d orbitals overlap.
• Metal atoms DO NOT lose their valence
  electrons in metallic bonding, rather they
  release them into a “Sea of Electrons”
• Although the atoms are bonded together,
  they are not bonded to any one particular
  atom, it is more like a large network.
Sea of Valence Electrons…
      Bonding Between Metals
• Results in an interaction that hold metal atoms
  together, however, it is not called a compound.
• Special Properties result from this interaction:
  – Malleable- pounded/ rolled into sheets (aluminum
  – Ductile- Drawn into wire. (copper wires)
  – Conductivity- the flow of electrons
  – Luster- Shiny-The narrow range of energy
    differences between orbitals allows electrons to be
    easily excited, and emit light upon returning to a
    lower energy level
          Metallic Properties
• Metals are good conductors of heat and light
• Metals have luster (shiny)
• The narrow range of energy differences between
  orbitals allows electrons to be easily excited, and
  emit light upon returning to a lower energy level
• Metals are Malleable- can be hammered into
  thin sheets
• Metals are ductile- ability to be drawn into wire
• Metallic bonding is the same in all directions, so
  metals tend not to be brittle
          Metallic Bond Strength
• Heat of Vaporization
• The ease with which atoms in a metallic solid can be
  separated from one another into individual gaseous
  atoms is related to bond strength
Chapter 6

6.5 Molecular
         VSEPR Model
   (Valence Shell Electron Pair Repulsion)

The structure around a given atom is
determined principally by minimizing electron
pair repulsions.
The model for predicting molecular shapes
VSEPR Theory: Repulsion between the sets of
valence-level electrons surrounding an atom
causes these sets to be oriented as far apart as
    VSEPR and Unshared Electron
• Unshared pairs take up positions in the geometry of
  molecules just as atoms do
• Unshared pairs have a relatively greater effect on
  geometry than do atoms
• Lone (unshared) electron pairs require more room
  than bonding pairs (they have greater repulsive
  forces) and tend to compress the angles between
  bonding pairs
• Lone pairs do not cause distortion when bond
  angles are 120° or greater
Predicting a VSEPR Structure

 Draw Lewis structure of each atom.
 Put pairs as far apart as possible.
 Determine positions of atoms from the way
electron pairs are shared.
 Determine the name of molecular structure
from positions of the atoms.
    Molecular Shapes: Linear
• When only 2 atoms are connected, the
  only possible shape is a straight line,
• Ex: Diatomics

                   Br      Br
  Molecular Shapes: Lone Pairs
• When lone pairs of electrons are present
  on the central atom, they may shift the
  shape of the molecule. These are pairs of
  electrons that are not involved in a bond.

            H       O       H

• They occupy space and provide negative
  repulsion forces against other electrons.
  They can repel more strongly than a
  bonded pair.
  Molecular Shapes: Repulsion
• Electrons want to spread out around the
  central atom.
• They want to be as far away from each
  other as possible, due to negative-
  negative repulsion.
• They want to maximize their distance.

                e- e-
        Molecular Shapes: Bent
• In water, the shape formed
  in called bent, which is a
  variation of a tetrahedron. A          O
  tetrahedral molecule has 1
  central atom and 4 attached.    H              H
  The bent molecule has 1
  central atom and 2 attached.
• The other 2 places are left “vacant” for the 2 lone
  pairs of electrons on oxygen, which occupy space
  and cause repulsion on the 2 bonded hydrogen
  atoms & cause them to be pushed downward, or
     Molecular Shapes: Linear
• Not all molecules with 3 atoms will have
  the bent shape.
• Carbon Dioxide has 2 double bonds; one
  to each oxygen. This gives CO2 a linear
  shape by placing the double bonds as far
  apart as possible.

            O       C      O
    Molecular Shapes: Pyramidal
• When you have 1 central atom with 3 attached,
  there are 2 possible arrangements.
• In ammonia, NH3, there are 3 hydrogen atoms
  bonded to a central nitrogen. Due to the lone
  pair on nitrogen, the hydrogen atoms are
  pressed downward to form the pyramidal
               H             H

 Molecular Shapes: Trigonal Planar
• The other possible arrangement for 1 central atom
  with 3 attached is called trigonal planar.
• In BCl3, the 3 chlorine atoms are bonded to a
  central boron atom. However, the boron atom does
  NOT have a lone pair. Therefore, there is not
  additional pressure placed on the 3 chlorine bonds,
  so they spread out equally around the central
  atom.                      Cl


                    Cl           Cl
   Molecular Shapes: Tetrahedral
• When 4 atoms are bonded to 1 central atom,
  the shape is called tetrahedral.
• In methane, CH4, 4 hydrogen atoms are
  bonded to 1 central carbon.

               H             H

      Table of Molecular Shapes
 # of                   Bond
          Structures            Shape Name
Atoms                  Angles

  2                     180°      Linear

  3                     105°       Bent

  4                     120°

  4                     107°     Pyramidal

  5                    109.5°   Tetrahedral
Table – VSEPR Structures
VSEPR & The Water Molecule
VSEPR & The Ammonia Molecule
VSEPR & Xenon Tetrafluoride

         Which one will it be???
VSEPR & Phosphorus Hexachloride
   How to Determine Polarity of a
• Draw the molecule’s shape according to the
  Lewis Structure and VSEPR
• Fill in the EN values, and determine the
  difference for each bond
• Determine the bond type for each bond in the
• Add in partial (δ) + and partial (δ) – symbols
• Determine if there is a divisible axis to
  separate the partial + from the partial -
              Molecular Polarity
• If your molecule has polar bonds, it may be a polar
• Polar molecules have special properties that result
  from partial positive and negative ends of the entire
• Polar molecules can also be called dipoles.
• These molecules are attracted to one another.

                                            δ = partial
            Molecular Polarity
• Water is a bent molecule,         2 δ-
  with 2 polar bonds.
• Since all of the negative
  charge is distributed on 1    H          H
  side of the molecule, and    δ+              δ+
  all the positive charge is
  on the opposite side,
  water is a polar molecule.
             Molecular Polarity
• Not all atoms with polar bonds are polar
• Carbon Dioxide has all polar bonds. However,
  its geometry prohibits overall polarity since
  there are no distinct ends splitting the positive &
  negative charges.
                δ-     2δ+       δ-
                O       C        O

• CO2, therefore, is a non-polar molecule.
The Blending of Orbitals
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent

   Lets look at a
   molecule of methane,

Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms
covalently bonded around it.
 Carbon ground state configuration
  What is the expected orbital notation of carbon
  in its ground state?

         Can you see a problem with this?

(Hint: How many unpaired electrons does this carbon
atom have available for bonding?)
            Carbon’s Bonding Problem
You should conclude that
carbon only has TWO
electrons available for
That is not enough for a full

How does carbon overcome this problem so that
it may form four bonds?
        Carbon’s Empty Orbital

The first thought that
chemists had was that
carbon promotes one of its
2s electrons…
                             …to the empty 2p
However, they quickly recognized a problem with such
an arrangement…

Three of the carbon-hydrogen bonds would involve an
electron pair in which the carbon electron was a 2p,
matched with the lone 1s electron from a hydrogen
This would mean that three of the bonds in a
methane molecule would be identical, because they
would involve electron pairs of equal energy.

         But what about the fourth bond…?

            Unequal bond energy
The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.

   Such a bond would have slightly less energy
   than the other bonds in a methane molecule.

              Unequal bond energy #2
This bond would be slightly different
in character than the other three
bonds in methane.

This difference would be measurable to a chemist
by determining the bond length and bond energy.

           But is this what they observe?

              Unequal bond energy #3
The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.

Chemists have proposed an explanation called

Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.
In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.

 These new orbitals have slightly MORE energy than
 the 2s orbital…
  … and slightly LESS energy than the 2p orbitals.

                 sp3 Hybrid Orbitals
               sp3    Hybrid Orbitals

Here is another way
to look at the sp3
and energy profile…
              sp Hybrid Orbitals
   While sp3 is the hybridization observed in methane,
   there are other types of hybridization that atoms

These include sp
hybridization, in
which one s
orbital combines
with a single p

Notice that this produces two hybrid orbitals, while
leaving two normal p orbitals
           sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals
from a p sublevel with one orbital from an s sublevel.

                                  Notice that one p
                                  orbital remains
Predicting the Geometry of
   Hybridized Orbitals
       Intermolecular Forces
• Forces of attraction between molecules
• Generally weaker than bonds that join atoms in
• Boiling point gives a rough estimate of
  intermolecular forces
• high bp = large attractive forces
• low bp = small attractive forces
 Molecular Polarity and Dipole-
        Dipole Forces
• Dipole- Created by equal but opposite charges
  that are separated by a short distance
• A dipole is represented by an arrow with a head
  pointing toward the negative pole and a crossed
  tail situated at the positive pole
  A molecule, such as HF, that has a center of
  positive charge and a center of negative
  charge is said to be polar, or to have a dipole

                   H F
                   +     
This symbol
means “partial”
      Polar Covalent Bonding

• δ = partial
         Dipole-Dipole forces
• The negative region of one molecule is attracted
  to the positive region of another molecule
• A polar molecule can induce a dipole in a
  nonpolar molecule by temporarily attracting its
Attraction between
oppositely charged
regions of neighboring
The Water Dipole
The Ammonia Dipole
• Water is a polar molecule.
• (A molecule with partially
  positive & negative ends due to
  differences in electronegativity)
• Its electrons are shared
  unequally so it is a dipole.
• Causes strong intermolecular
           Hydrogen Bonding
• The intermolecular force in which a hydrogen
  atom that is bonded to a highly electronegative
  atom (F, O, N) is attracted to an unshared pair of
  electrons of an electronegative atom in a nearby
• It is usually represented by dotted lines
Hydrogen Bonds
• When the partial positive, or
  hydrogen end of a molecule is
  attracted to the partial negative,
  or oxygen end of another
  molecule we call it a Hydrogen
• It is an attraction!
• Hydrogen, and Fluorine, Oxygen
  & Nitrogen. (F-O-N)
  (high electronegativity values)

• Diagram of H-bonds
• Solid lines are normal bonds
• Dotted lines are H-bonds
• Molecules containing
  hydrogen bonds, like
  water, have very high
• Why might this occur?
• Hydrogen bonds are
  strong -holds molecules
• It takes a lot of energy to break apart these
  attractions to liberate each molecule into the
  gaseous state.
 Hydrogen Bonding
Bonding between
hydrogen and more
neighboring atoms such
as oxygen and nitrogen

Hydrogen bonding is
used in Kevlar, a
strong polymer used in
bullet-proof vests.
in Water
Hydrogen Bonding Between
    Ammonia & Water
           London Dispersion Forces
                 aka Van Der Waals Forces
                   The temporary separations of charge
                   that lead to the London force
                   attractions are what attract one
                   nonpolar molecule to its neighbors.
                  •All molecules experience London
  Fritz London    •London forces increase with the size
   1900-1954      of the molecules.

• London forces are the only forces of attraction
 among noble-gas atoms, nonpolar, and slightly
 polar molecules
London Forces in Hydrocarbons
Comparison of Boiling Points &
        Bond Types
Relative Magnitudes of Forces
 The types of bonding forces vary in their
 strength as measured by average bond
Strongest    Covalent bonds (400 kcal)

             Hydrogen bonding (12-16 kcal )

             Dipole-dipole interactions (2-0.5 kcal)

Weakest      London forces (less than 1 kcal)

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