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Organic Chemistry, Exam 1 Review

by: Hannah Malcomson

Organic Chemistry, Exam 1 Review CHEM141A

Marketplace > University of Vermont > Chemistry > CHEM141A > Organic Chemistry Exam 1 Review
Hannah Malcomson

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About this Document

This review sheet covers the last four weeks of class, but does not include what will be taught tomorrow in class. I will upload tomorrow notes a few hours after class ends.
Organic Chemistry 1
Study Guide
Organic Chemistry, Organic Chem
50 ?




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This 4 page Study Guide was uploaded by Hannah Malcomson on Sunday September 25, 2016. The Study Guide belongs to CHEM141A at University of Vermont taught by Wurthmann in Fall 2016. Since its upload, it has received 8 views. For similar materials see Organic Chemistry 1 in Chemistry at University of Vermont.


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Date Created: 09/25/16
Exam 1 Review Sheet: Definitions:  Organic molecule- a compound containing carbon atoms  Constitutional isomers- compounds with the same molecular formula, but have a different structure (atoms are connected in different orders)  Covalent bond- bond formed by sharing of electrons  Lewis structures: drawings to convey the electron interactions between molecules  Octet rule- every atom (with a few exceptions), will be likely to form bonds in order to have 8 electrons in its valence shell  Formal charge- a charge associated with any atom that does not exhibit the appropriate number of valence electrons  Types of bonds o Covalent- different in electronegativity is <0.5. formed by an overlap of atomic orbitals. o Polar covalent- difference in electronegativity is between 0.5 and 1.7 o Ionic- difference in electronegativity is >1.7. not actually bonds- forces of attraction hold atoms together  Molecular orbital theory- describes a bond as the constructive interference between two overlapping atomic orbitals.  Bond strength: Single (alkane) < Double (alkene) < Triple (alkyne)  Steric number: sum of all sigma bonds and lone pairs  Dipole moments- the center of negative charge and positive charge are separated from one another by a certain distance. Indicative of polarity  Intermolecular forces- attractive forces between molecules. 3 types: o Dipole- dipole o Fleeting dipole- dipole: London dispersion. The temporary positive and negative charges may no be equal, and a temporary, weak attraction will form o Hydrogen bonding- the difference in electronegativies between hydrogen and the atom it is bonded to will lead to the creation of a dipole, which will be very strong since hydrogen is a very small atom, and the charges will be very close together.  Hydrophobic- not soluble in water  Hydrophilic- soluble in water  Like dissolves like- polar compounds will be soluble in polar solvents, and vice versa  Hybridized orbitals: 3 o sp hybridized- tetrahedral. 3 p- orbitals are mixed with s- orbital. All single bonds. o sp hybridized- trigonal planar. One double bond. o sp- hybridized- one triple bond  Solubility- organic compounds and water o Ionic bonds- any compounds with a charge will be soluble in water, as water is a polar molecule.  Functional groups: a grouping of atoms, bonded together, with a predictable reactivity/ chemical behavior. A molecule will often have multiple functional groups.  Functional groups with all single bonds: o Alcohols: all single bonds between carbons atoms, with an –OH group off one of the carbons o Ether: all single bonds, with an oxygen atom between two carbon atoms o Amine: all single bonds with a nitrogen atom attached to a carbon atom  Functional groups with double bonds o Aldehyde: a double bond between a carbon and an oxygen, at the end of the chain o Ketone: a double bond between a carbon and an oxygen, in the middle of the chain o Carboxylic acid; a double bond between a carbon and an oxygen, with a C-OH off that same carbon o Ester: a double bond between a carbon and an oxygen, with a C- O-R coming off that same carbon atom o Amide: a double bond between a carbon and an oxygen, with a C-N-R coming off that same carbon atom  Curved arrows do the following o Form a bond between two atoms by moving electrons to the space between the atoms o Break a bond between two atoms by pulling electrons from the space between the atoms o Generate charges by moving electrons onto an atom  Electrons come from 1 of 3 places to create or break a bond: o Negatively charged atoms o Lone pairs o Pi- bonds (double or triple bonds)  Resonance: explains movement of electrons within a structure in reference to the structures stability o If electrons (charges) can be shared by multiple atoms, the electrons are more stable  Multiple stable resonance structures make a resonance hybrid, which is more stable because the positive and negative charges are spread over multiple atoms, meaning there is no entirely positive or entirely negative charge on any atom at one time.  Patterns to look for when predicting resonance and using curved arrows: o Allylic pattern  o Neighboring atoms of differing electronegativity   Guidelines: o Curved arrows move electrons o For resonance, avoid breaking single bonds, as breaking single bonds indicates a reaction, not resonance o Never exceed the octet rule for second row elements o Minimize charges. Bonds stabilize atoms. Choose structures with more bonds vs. more charges o Electronegative atoms (N, O, Cl, Br) are best for hosting negative charges (may even host positive charges)  Bronstead- Lowry definitions o Acid= H+ donor (proton donor) o Base= H+ acceptor  Lewis definitions: o Lewis Acid: electron pair acceptor o Lewis Base: electron pair donor (start of curved arrows)  Acids and Bases: o pKa= relative acidity of a compound  Used to predict the favored side of an equilibrium  Since the more acidic molecule in a reaction is more likely to donate an H+, it is less stable. The weaker acid will be more stable, and is therefore the favored side of the equilibrium.  Larger pKa= weaker acid= favored side of the equilibrium o Process for Assessment: (use pKa’s to find favored side of equilibrium)  Identify the base (H+ acceptor, frequently negative charge or lone pairs)  Identify acidic hydrogens (label ∂+ and ∂-)  Draw curved arrows  Draw structures of conjugate acids and conjugate bases based on the mechanism  Assign pKa values, and predict favored side. Higher pKa value= favored side o What to do without pKa values? Estimate acid/ base strength  Acidity increases across the row of the P.T.  Down a column of the Periodic table, acidity increases. Larger atom- more effective distribution of electrons= more stable  Resonance- sharing of electrons (negative charge) over multiple atoms stabilizes charge.  Induction- electronegative atoms create a polar covalent bond which can deplete another atom of it’s electron density. Depleting a negative charge stabilizes a molecule.  Hybridization- the negative charge in an sp- hybrid orbital is more spherical compared to an sp or sp –hybrid orbital. This pulls the electrons closer to the nucleus, stabilizing the negative charge.


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