Molecular Orbital Theory - A Brief Review

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Molecular Orbital Theory- A Brief Review
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What is an Atomic Orbital?

• Heisenberg Uncertainty Principle states that it
  is impossible to define what time and where an
  electron is and where is it going next. This
  makes it impossible to know exactly where an
  electron is traveling in an atom.
• Since it is impossible to know where an electron
  is at a certain time, a series of calculations
  are used to approximate the volume and time in
  which the electron can be located. These regions
  are called Atomic Orbitals. These are also known
  as the quantum states of the electrons.
• Only two electrons can occupy one orbital and
  they must have different spin states, ½ spin and
  ½ spin (easily visualized as opposite spin
  states).

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s and p Atomic Orbitals

• These are some examples of atomic orbitals
• s orbital (Spherical shape) There is one S
  orbital in an s subshell. The electrons can be
  located anywhere within the sphere centered at
  the atoms nucleus.

• p Orbitals (Shaped like two balloons tied
  together) There are 3 orbitals in a p subshell
  that are denoted as px, py, and pz orbitals.
  These are higher in energy than the corresponding
  s orbitals.

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Electron Configuration

• Every element is different.
• The number of protons determines the identity of
  the element.
• All chemistry is done at the electronic level
  (that is why electrons are very important).
• Electronic configuration is the arrangement of
  electrons in an atom. These electrons fill the
  atomic orbitals

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Electron Configuration of Li

• The arrows indicate the value of the magnetic
  spin (ms) quantum number (up for 1/2 and down
  for -1/2)
• The occupancy of the orbitals would be written in
  the following way
• 1s22s1

http//wine1.sb.fsu.edu/chm1045/notes/Struct/EConf
ig/Struct08.htm
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Electron Configurations and Box Diagrams
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What are Valence Electrons?

• The valence electrons are the electrons in the
  last shell or energy level of an atom.

The lowest level (K), can contain 2
electrons. The next level (L) can contain 8
electrons. The next level (M) can contain 8
electrons.
Carbon - 1s22s22p2  - four valence electrons
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Valence Bond Theory

• Explains the structures of covalently bonded
  molecules
• how bonding occurs
• Principles of VB Theory
• Bonds form from overlapping atomic orbitals and
  electron pairs are shared between two atoms
• A new set of hybridized orbitals can form
• Lone pairs of electrons are localized on one atom

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Molecular Orbital (MO) Theory

• Explains the distributions and energy of
  electrons in molecules
• Useful for describing properties of compounds
• Bond energies, electron cloud distribution, and
  magnetic properties
• Basic principles of MO Theory
• Atomic orbitals combine to form molecular
  orbitals
• Molecular orbitals have different energies
  depending on type of overlap
• Bonding orbitals (lower energy than corresponding
  AO)
• Nonbonding orbitals (same energy as corresponding
  AO)
• Antibonding orbitals (higher energy than
  corresponding AO)

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Formation of Molecular Orbitals

• Recall than an electron in an atomic orbital can
  be described as a wave function utilizing the
  Schröndinger equation. The waves have positive
  and negative phases. To form molecular orbitals,
  the wave functions of the atomic orbitals
  combine. How the phases or signs combine
  determine the energy and type of molecular
  orbital.

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Formation of Molecular Orbitals

• Bonding orbital the wavefuntions are in-phase
  and overlap constructively (they add).
• Bonding orbitals are lower in energy than AOs
• Antibonding orbital the wavefunctions are
  out-of-phase and overlap destructively (they
  subtract)
• Antibonding orbitals are higher in energy than
  the AOs
• When two atomic orbitals combine, one bonding and
  one antibonding MO is formed.

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Rules for Filling Electrons in Molecular Orbitals

• Electrons go into the lowest energy orbital
  available to form lowest potential energy for the
  molecule.
• The maximum number of electrons in each molecular
  orbital is two. (Pauli exclusion principle)
• One electron goes into orbitals of equal energy,
  with parallel spin, before they begin to pair up.
  (Hund's Rule.)

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Molecular Orbital Diagram

• In atoms, electrons occupy atomic orbitals, but
  in molecules they occupy similar molecular
  orbitals which surround the molecule.
• The two 1s atomic orbitals combine to form two
  molecular orbitals, one bonding (s) and one
  antibonding (s).
• Each line in the diagram represents an orbital.
• The electrons fill the molecular orbitals of
  molecules like electrons fill atomic orbitals in
  atoms

• This is an illustration of molecular orbital
  diagram of H2.

• Notice that one electron from each atom is being
  shared to form a covalent bond.

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Molecular Orbital Diagram (H2)
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Energy Diagram for Sigma Bond Formation by
Orbital Overlap
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Examples of Sigma Bond Formation
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Overlap of 2px Orbitals
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Overlap of 2py 2pzOrbitals
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MO Diagram for O2
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Bond Order and Bond Stability
A bond order equal to zero indicates that there
are the same number of electron in bonding and
antibonding orbitals The greater the bond order,
the more stable the molecule or ion. Also, the
greater the bond order, the shorter the bond
length and the greater the bond energy.
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Heteronuclear Diatomic Molecules

• Molecular orbital diagrams for heteronuclear
  molecules have skewed energies for the combining
  atomic orbitals to take into account the
  differing electronegativities.
• The more electronegative elements are lower in
  energy than those of the less electronegative
  element.

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Energy Level Diagram for NO
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The Energy Level Diagram for HF
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Molecular Orbital Diagram (CH4)
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Molecular Orbital Diagram (H2O)