Bonding Theories

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Bonding Theories
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Hybrid Orbitals

• VSEPR does not describe type of bonds
• ORBITAL HYBRIDIZATION provides information
  about molecular bonding and molecular shape
• HYBRIDIZATION several atomic orbitals mix to
  form the same total number of equivalent hybrid
  orbitals

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Consider Methane CH4
1s
2s
2p
H
Due to the promotion of an electron towards
hybridization, there is a greater overlap which
results in an unusually strong single covalent
bond
H
H
H
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Bond Types

• Covalent bonds differ in terms of how the bonded
  atoms share the electrons.
• The character of the bonds in a given molecule
  depends on the kind and number of atoms joined
  together.

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Three types of Bonds

• Nonpolar covalent
• Polar covalent
• Ionic

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Nonpolar Covalent Bonds

• The bonding electrons are shared equally between
  the nuclei of the atoms sharing electrons
• Examples diatomic molecules
• The electronegativity difference between the
  atoms is very small

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Polar Covalent Bonds (Polar Bonds)

• A covalent bond between atoms in which the e- are
  shared unequally
• The more EN atom attracts more strongly and gains
  a slight (less than 1) negative charge.

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• The less EN atom has a slight positive charge.
• Example Water
• Oxygen has a much higher EN value than hydrogen
• Oxygen will have a slight negative charge,
  hydrogen will have a slight positive one

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THE REAL STORY

• Three basic bond types are defined by the
  DIFFERENCE in electronegativity (?EN) between the
  atoms
• Ionic 1.7
• Polar Covalent 0.6 1.7 (unequal sharing)
• Non-polar Covalent 0.6

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Increasing ?EN
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• To find the DIFFERENCE subtract the lower value
  from the higher value.
• ?EN (larger EN)- (smaller EN)

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Examples
Compound F2 HF LiF
?EN 4.0 - 4.0 0 4.0 - 2.1 1.9 4.0 - 1.0 3.0
Type of Bond Nonpolar covalent Polar covalent Ionic (non-covalent)
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• Generally, the element closest to fluorine will
  be relatively negative, and the element furthest
  from fluorine will be relatively positive
• Polarity is indicated by

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• Indicate the positive and negative ends of each
  of the following bonds by using the symbols
• a. S-O
• b. C-N
• c. S-P
• d. C-F
• e. H-Br

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• A. -
• B. -
• C. -
• D. -
• E. -

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• Geometry or shape of a molecule can be extended
  to larger molecules
• Helpful in determining properties of compound

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VSEPR THEORY

• Valence Shell Electron Pair RepulsionTheory
• Repulsion between electron pairs causes molecular
  shapes to adjust so that valence electron pairs
  stay as far apart as possible
• Unshared pairs of e- are also important in
  predicting shapes of molecules

• Check out the site below to find out more about
  molecular geometry.
• http//intro.chem.okstate.edu/1314F00/Lecture/Chap
  ter10/VSEPR.html
• Remember predicted shapes may be different from
  actual shapes.
• Draw Lewis Structures First

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Polar Molecules

• One end of the molecule is slightly negative, the
  other end is slightly positive.
• A molecule with two poles is called a dipolar
  molecule.
• The effect of polar bonds on the polarity of an
  entire molecule depends on the shape of the
  molecule and the orientation of the polar bonds.

Lets see why on the next page
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Four General Rules for molecular polarity

• 1. a molecule that is symmetrical is nonpolar,
  it does not matter how polar the individual bonds
  are (if all outer atoms are the same)
• 2. a nonsymmetrical molecule is polar if the
  bonds are polar

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• 3. a molecule with more that one type of atom
  attached to the central atom is often
  nonsymmetrical and therefore polar
• 4. A central atom with nonbonding electron pairs
  is often nonsymmetrical and polar

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Polar Molecules continued

• Carbon dioxide has two polar bonds
• It is linear
• Bond polarities are on the same axis in opposite
  directions and will canceltherefore it is
  nonpolar

• Water has two polar bonds
• It is bent
• Bond polarities do not cancel, making this a
  polar molecule

Remember we are referring to the polarity of
the entire molecule here, not just the individual
bonds within the molecule
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• Ex construct the Lewis structure and predict the
  polarity of each of the following
• A. CH3Cl
• B. CS2
• C. PH3
• D. SiF4
• E. CO32-

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• A. Polar with Cl as the negative end
• B. Nonpolar
• C. Polar with P at the negative end
• D. Nonpolar
• E. Nonpolar because of resonance, charge on the
  ion is distributed evenly over the ion charged
  but not polar

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Do Now

• 1. If two atoms have a large difference in
  electronegativity, is the bond between them
  polar, nonpolar or ionic?

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• 2. Which bonds in the following pairs have less
  ionic character?
• Na-Cl or Ca-Cl
• Cs-Cl or Ba-Cl
• Fe-I or Fe-F

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• 3. For each of the following bonds, draw a
  figure indicating the direction of the bond
  dipole, including which end is positive and which
  is negative
• H-C
• N-O
• N-S

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Intermolecular attractions

• intermolecular attractions are weaker than either
  ionic or covalent bonds.
• They are, however, still important and can
  determine things like whether a molecular
  compound is a solid, liquid, or gas.

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IntERmolecular
Intermolecular force
Intramolecular force
Intermolecular forces weaker than intramolecular
force
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• Halogens change from gas, to liquid, to solid
  because of these forces
• Responsible for physical and chemical properties
  of a substance freezing point and boiling point

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• use the force...

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Van der Waals Forces

• Types of Van der Waals forces
• Dipole interactions - occur when polar molecules
  are attracted to one another They are similar to
  but weaker than ionic bonds
• Dispersion forces weakest of the molecular
  interactions

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Dipole interactions

• Ion-Dipole- attractive forces between an ion and
  a polar molecule

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Ion- dipole
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Dipole- Dipole

• Dipole-Dipole interactions- attractive forces
  between the negative end of one POLAR molecule
  and the positive end of another POLAR molecule

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Dipole - Dipole
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Hydrogen Bonds-type of dipole - dipole

• Attractive forces in which a hydrogen covalently
  bonded to a very electronegative atom (F, O, N)
  is also weakly bonded to an unshared electron
  pair of another electronegative atom
• 5x stronger than average dipole-dipole

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Hydrogen bond
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Hydrogen bonding

• Strongest intermolecular forces- why?
• - Always involves hydrogen- 1 e- and 1 p
• Hydrogens valence e- are not shielded from the
  nucleus by a layer of underlying e-
• F,O,N have large electronegativity lone pairs
  on these increase negative charge

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Trends in boiling points
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Trends in boiling points

• As you move down a group the molar mass increases
  so does boiling point
• Takes more energy to push a large atom into the
  gas phase
• water and ammonia- have high bp because only
  molecules on the graph that have hydrogen bonding
  
• More energy to break H-bonds and send into gas
  phase

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Ion induced dipole

• An ion can induce or create a dipole in a
  nonpolar species
• Electrons are attracted or repel the nearby ion
• Causes temporary increase in electron density on
  one side of the molecule
• Then electrons re-align
• Attraction constantly being turned on and off

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Ion induced dipole
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Hemoglobin

• Globular protein
• 4 segments in
• tetrahedral like arrangement
• 4 segments of protein structures - fold to form
  pocket contain non protein heme group
• Hydrogen bonds stabilize the helical sections
  inside this protein, causing attractions within
  the molecule, folding each polypeptide chain into
  a specific shape

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Heme

• Hemogobin
• Fe2 induces a dipole in the nonpolar O2
• The iron ion, which is the site of oxygen
  binding, coordinates with the fournitrogens in
  the center of the ring, which all lie in one
  plane.
• The iron is bound strongly (covalently)
• A sixth position can

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• reversibly bind oxygen by a coordinate covalent
  bond
• , completing the octahedral group of six ligands.
• Oxygen binds in an "end-o
• n bent" geometry where one oxygen atom binds Fe
  and the other protrudes at an angle.

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• reversibly bind oxygen by a coordinate covalent
  bond
• completing the octahedral group of six ligands.
• Oxygen binds in an "end-on bent" geometry where
  one oxygen atom binds Fe and the other protrudes
  at an angle.
• octahedron.

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• Dispersion Forces are the weakest of all
  molecular interactions.
• As electrons move, they repel other electrons and
  cause them to move (a chain reaction of sorts)
• The force increases as the of e- increases

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• www.youtube.com/watch?vLGwyBeuVjhUfeaturerelate
  dhttp//www.youtube.com/watch?vLGwyBeuVjhUfeatur
  erelated
• dispersion forces

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• http//www.youtube.com/watch?voSwlZ4Sio9Qfeature
  related
• Water video

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Molecular OrbitalsWhen two atoms combine, their
atomic orbitals overlap and form molecular
orbitals

• Sigma Bonds s
• When two atomic orbitals overlap side by side to
  form a molecular orbital a sigma bond is formed
• Can be overlap of two s orbitals or one s and one
  p

• Pi Bonds p
• Orbitals overlap side by side
• Electron density above and below the nuclei

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• Single bond sigma bond
• Double bond one sigma and one pi bond
• Triple bond one sigma and two pi

• Hybridization

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Hybrid orbital model

• Several atomic orbitals mix to form the same
  number of equivalent hybrid orbitals
• hybridization makes the VSEPR work mathematically

• orbital formation
• sp, sp2, and sp3
• hybrid orbital- a set of orbitals formed from the
  combination of two or more atomic orbitals with
  different energies

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The End