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Organic Chemistry, CHMY 321, Fall 2016

by: Adriana

Organic Chemistry, CHMY 321, Fall 2016 CHMY 321-001

Marketplace > Montana State University - Bozeman > CHMY 321-001 > Organic Chemistry CHMY 321 Fall 2016

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About this Document

Covers the first weeks of notes. Is meant as a supplement, not as a replacement for notes taken in class. Is more of a brief overview, since the first week was mostly review.
Organic Chemistry I
Holmgren, Steven
Class Notes
Organic Chemistry, Chemistry, MSU-Bozeman




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This 4 page Class Notes was uploaded by Adriana on Sunday September 4, 2016. The Class Notes belongs to CHMY 321-001 at Montana State University - Bozeman taught by Holmgren, Steven in Fall 2016. Since its upload, it has received 115 views.


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Date Created: 09/04/16
OrganicChemistry I-Fall2016 Review of chemistry: The Atom: - Made of:  Electrons (-)  Protons (+)  Neutrons (neutral) Electron Configuration - Find how many electrons are in the element - Use the Aufbau Principle  fill the lowest energy orbitals first - Each orbital has a max of 2 electrons (Pauli Exclusion Principle), these must have opposite spin - Degenerate orbitals are filled all with 1 electrons of the same spin, before being paired up (Hund’s Rule) - Count the amount of valence electrons, these are the bonding electrons  Example: Boron  !s 2s 2p , has 3 valence electrons from the outer 2 shells. 2p __ __ __ 2s __ 1s __ Bonding - Octet Rule  Atoms want to have full shells = stability  8 electrons in shell  Exceptions to the Rule: + Incomplete octets  Boron only bonds with 6 electrons + Expanded Octets  Elements in third row and beyond can have up to 12-14 electrons in a shell Electronegativity - How much electrons are attracted to an atom - Increases from left to right in periodic table, and from bottom to top - Noble gases don’t react (usually) since they are stable Ionic bonding - Transfer of electrons to create a cation and anion - Electrostatic attraction - Particularly groups I & II with N,O,F,Cl,Br,F,S - Example: Na + Cl  NaCl Covalent Bonding - Electron sharing between two atoms - Equal sharing does not always occur - Example: Cl-Cl Geometry (VSEPR) - Valence Shell Electron Pain Repulsion theory - Steric #= # of lone pairs + # of sigma bonds  Shows the number of bonded and non-bonded electrons pairs that are repelling one another.  Treat pi bonds as if they are not there in this calculation - These repulsions cause the electrons to arrange themselves with maximal distance between each other, creating 3-D structures. - Figure 1.43 on page 46 in the book shows predicting geometry. Also, there is table 6.5 in Tro, 2015, Chemistry: Structure and Properties on page 216, which gives examples along with the geometry and other such information. Non-polar vs. polar bonds Electronegativity Bond type Example difference < 0.5 Covalent H-H 0.5-1.7 Polar Covalent O-H > 1.7 Ionic NaCl Molecular Polarity: - Depends on molecular shape and if there are polar bonds present - Dipole moments- used to indicate polarity - Use sum of atomic vectors to get overall molecular dipole moment (shown as a vector) (see figure 1.45 pg. 49 in book for example) - Measured in Debye (D) Intermolecular interactions: - London dispersion forces (Van der Waals forces)  Fleeting dipole-dipole interactions  Example: Hydrocarbons, pg. 55 - Hydrogen Bonding  Is a specific kind of dipole-dipole interactions  Hydrogen (H) is connected to an electronegative atom, such as O, meaning H is slightly positive, therefore it interact with a lone pair of electrons from another atoms, which are slightly negative.  Happens with protic compounds  compound that have a proton connected to an electronegative atom.  Is a strong interaction  Example: pg. 52 figure 1.49 Molecular Orbital Theory - Bonds are the constructive interference of two overlapping atomic orbitals that make combine to make a molecular orbital that consists of bonding (constructive interference of atomic orbitals) and antibonding (destructive interference) molecular orbitals.  Antibonding Molecular Orbitals  higher energy than bonding Molecular Orbitals  Example: figure 1.15 pg 35  Bond order = (# of bonding electrons - # of antibonding electrons)/2 Hybridization - In order to find the hybridization of an atom in a compound can find the steric number of each.  Skill builder 1.8 is a good example of this, pg. 46-47  Steric numbers of 4 are sp3 2  Steric numbers of 3 are sp  Steric numbers of 2 are sp - Can also look at it from the perspective of electron geometry (Table in Tro, 2015 pg. 248 table 7.1)  If there are 3 electron groups  linear configuration sp hybridization 2  If there are 3 electron groups  Trigonal Planar configuration  sp hybridization  If there are 4 electron groups  Tetrahedral configuration  sp hybridization Lewis dot structure - Set up:  Find how many valence electrons there are  Connect atoms that make 1+ bond(s)  Add Hydrogen bonds  Add the rest of the electron pairs  Double check the octet rule and find formal charge = (valence electrons – (lone pairs + ½ bonded electrons))  Find partial charge + I am assuming you know the basics of this, but if you need help look at page 71. Condensed structural formula - Set up:  Central atom first  Then any H or attached atoms  Then the next central atom  Then the attached atoms to that  Repeat  Lone pairs are not included + Examples: pg. 71 Bond-line structural formula - Bonds are shown as lines - Carbons are assumed to be where lines bend - Hydrogens are assumed unless it is a heteroatom - Other elements are written in  Example: See page 75 3-D structure (Stereochemistry) - Shows structure in 3-D form - Like in VSPER theory, but with bond-line format (as shown in class) First 8 functional groups (need to be memorized by first exam) - Alkyl Halide  A carbon to halogen bond  C-X, X being F, Cl, Br, or I - Alkene  A carbon double bonded to a carbon  C=C - Alkyne  A carbon triple bonded to a carbon  C= C - Alcohol  A carbon attached to an OH group  C-OH - Ether  Carbon bonded to Oxygen bonded to Carbon  C-O-C - Thiol  Carbon bonded to SH  C-SH - Sulfide  Carbon bonded to Sulfur bonded to Carbon  C-S-C - Arene (Aromatic)  A cyclic array of single and double bonds - See table 2.1 on pg. 78 for examples of functional groups.


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