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[CHEM 2E03] - Final Exam Guide - Ultimate 102 pages long Study Guide!

102 pages131 viewsWinter 2014

Department
Chemistry
Course Code
CHEM 2E03
Professor
Dr.S
Study Guide
Final

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McMaster
CHEM 2E03
FINAL EXAM
STUDY GUIDE
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Chapter 1: Review
1.2 Structural Theory
Constitutional isomers: same molecular formula but different physical properties and different names
Valence: number of bonds usually formed by each element
-Monovalent: forms one bond (eg. Hydrogen, halogens)
-Divalent: forms two bonds (eg. Oxygen)
-Trivalent: forms three bonds (eg. Nitrogen)
-Tetravalent: forms four bonds (eg. Carbon)
1.3 Electrons, Bonds, and Lewis Structures
-covalent bond: result of two atoms sharing a pair of electrons (eg. H-H)
- when a covalent bond forms, there is a decrease in energy indicated by a negative ΔH value
- electrons move together to minimize repulsive forces and maximize attractive forces
oprovides a net force of attraction which lowers the energy of the system
- valence electrons: the electrons in the outermost shell of an atom
o# of valence electrons is identified by its group # in the periodic table
1.4 Identifying Formal Charges
- a formal charge is associated with any atom that doesn’t exhibit the appropriate # of valence e’s
1. Determine the appropriate number of valence e’s for an atom
oGroup number indicates appropriate valence #
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2. Determine whether the atom exhibits the appropriate number
oEach bond represents 2 shared e’s  split each bond equally.. 1 e counts per bond
1.5 Induction and Polar Covalent Bonds
-electronegativity: measure of the ability of an atom to attract electrons
oEN difference < 0.5 : covalent bond
o0.5 < EN difference < 1.7 : polar covalent bond
Induction: Withdrawal of e’s towards one atom, shown with arrow w/ a cross
Formation of partial pos/neg charges, shown by δ+ / -
oEN difference > 1.7: ionic bond (e’s not shared)
1.7 Valence Bond Theory
- Like waves, when atomic orbitals overlap, they can interfere constructively or destructively
- A bond is the sharing of electron density btwn 2 atoms as a result of constructive interference of
their atomic orbitals
oEg. Overlap of 1s atomic orbitals of 2 H atoms
-Sigma bond: circular symmetry with respect to the bond axis  all single bonds are sigma bonds
1.8 Molecular Orbital Theory
- describes a bond in terms of the constructive interference between 2 overlapping orbitals
-linear combination of atomic orbitals (LCAO): atomic orbitals are mathematically combined to
produce new orbitals called molecular orbitals
-atomic orbital: region of space associated with an individual atom
-molecular orbital: associated with the entire molecule
omolecule considered a single entity held together by many electron clouds
- molecular orbitals are filled with electrons in a particular order like atomic orbitals
- electrons first occupy the lowest energy orbital with a max of 2 e’s per orbital
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