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by: Lauren Faris


Marketplace > University of Alabama - Tuscaloosa > Science > Ch 101 > CHEM 101 THIRD WEEK NOTES
Lauren Faris

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These are the cliff notes from the information we learned in chapter 7.
Chemistry 101 008
Dr. Bakker
Class Notes
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This 3 page Class Notes was uploaded by Lauren Faris on Friday October 14, 2016. The Class Notes belongs to Ch 101 at University of Alabama - Tuscaloosa taught by Dr. Bakker in Fall 2016. Since its upload, it has received 11 views. For similar materials see Chemistry 101 008 in Science at University of Alabama - Tuscaloosa.




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Date Created: 10/14/16
CHEM 101 THIRD WEEK NOTES By Lauren Faris Chapter 7 Cliff Notes  Oxygen is paramagnetic.  Hybridization is where you mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals.  Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process.  The valence electrons of the atoms in a molecule reside in quantum mechanical atomic orbitals.  A chemical bond results when these atomic orbitals interact and there is a total of two electrons in the new molecular orbital. The electrons must be spin paired.  According to valence bond theory, bonding takes place between atoms when their atomic or hybrid orbitals interact.  To interact, the orbitals must do either of the following: o Be aligned along the axis between the atoms. o Be parallel to each other and perpendicular to the interatomic axis.  Some atoms hybridize their orbitals to maximize bonding. More bonds = more full orbitals = more stability.  Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals. sp, sp2 , sp3 , sp3d, sp3d2  The same type of atom can have different types of hybridization.  A sigma (σ) bond results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei.  Either standard atomic orbitals or hybrids • s to s, p to p, hybrid to hybrid, s to hybrid, etc. • A pi (π) bond results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei. Between unhybridized parallel p orbitals.  The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore, σ bonds are stronger than π bonds.  Atom with two electron groups are linear shape with a 180° bond angle.  Valence Bond (VB) theory predicts many properties better than Lewis theory. – Bonding schemes, bond strengths, bond lengths, bond rigidity.  Resonance hybrids: VB theory presumes the electrons are localized in orbitals on the atoms in the molecule, so doesn’t really address resonance structures.  In MO theory you apply Schrödinger’s wave equation to the molecule to calculate a set of molecular orbitals. The equation solution is estimated. The estimated solution is evaluated and adjusted until the energy of the orbital is minimized.  The simplest guess starts with the atomic orbitals of the atoms adding together to make molecular orbitals; this is called the linear combination of atomic orbitals (LCAO) method.  Because the orbitals are wave functions, the waves can combine either constructively or destructively.  When the wave functions combine constructively, the resulting molecular orbital has lower energy than the original atomic orbitals. Called a bonding molecular orbital. o Designated: σ, π o Most of the electron density between the nuclei.  When the wave functions combine destructively, the resulting molecular orbital has higher energy than the original atomic orbital. Called an antibonding molecular orbital. o Designated: σ*, π* o Most of the electron density outside the nuclei. o Nodes between nuclei.  Use Aufbau approach for MO’s (as we did for individual atoms).  Electrons in bonding MOs are stabilizing. o Lower energy than the atomic orbitals.  Electrons in antibonding MOs are destabilizing. o Higher in energy than atomic orbitals. o Electron density located outside the internuclear axis. o Electrons in antibonding orbitals cancel stability gained by electrons in bonding orbitals.  Bond order = ½ (# Bonding Electrons – # Antibonding Electrons).  Bond order = difference between number of electrons in bonding and antibonding orbitals.  Molecular orbitals (MOs) are a linear combination of atomic orbitals (AOs).  MOs are named by type: σ, π, with a subscript to indicate what AOs they were formed from.  When the combining atomic orbitals are identical and of equal energy, the contribution of each atomic orbital to the molecular orbital is equal.  When the combining atomic orbitals are different types and energies, contributions to the MOs are different.  The more electronegative an atom is, the lower in energy are its orbitals. • Lower energy atomic orbitals contribute more to the bonding MOs.  Higher energy atomic orbitals contribute more to the antibonding MOs.  Nonbonding MOs remain localized on the atom donating its atomic orbitals.  When many atoms are combined together, the atomic orbitals of all the atoms are combined to make a set of molecular orbitals, which are delocalized over the entire molecule. This gives results that better match real molecule properties than either Lewis or valence bond theories.  Band Theory: o Electrons become mobile when they make a transition from the highest occupied molecular orbital into higher energy empty molecular orbitals. o These occupied molecular orbitals are referred to as the valence band. o The unoccupied orbitals the conduction band.


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