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by: Alice Hsu


Marketplace > Dartmouth College > Chemistry > Chemistry 51 > CHEM 51 MIDTERM 1 STUDY GUIDE
Alice Hsu

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this guide covers the first four chapters of our organic chemistry. we review acids and bases, structural theory, valence bonding, hybridization, and nomenclature.
Organic Chemistry
Peter Jacobi
Study Guide
Chemistry, Organic Chemistry, acids and bases, valence electrons, hybridization, nomenclature, Alkanes, cycloalkane conformations, Cycloalkanes, Conformations
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This 7 page Study Guide was uploaded by Alice Hsu on Saturday September 24, 2016. The Study Guide belongs to Chemistry 51 at Dartmouth College taught by Peter Jacobi in Fall 2016. Since its upload, it has received 5 views. For similar materials see Organic Chemistry in Chemistry at Dartmouth College.




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Date Created: 09/24/16
CHEM 51 MIDTERM 1 STUDY GUIDE *be able to provide examples for everything. BASICS OF STRUCTURAL THEORY: 1828: Fredrich Wohler discovers that organic compounds can be synthesized from inorganic compounds and gives birth to structural theory Organic compounds: contain a carbon backbone and hydrogens. Ionic bond: transfer of e- Covalent bond: sharing of e- Simplistic model of an atom: *mostly space *each shell contains a maximum number of e-: 2, 8, 18 *maximum stability associated with filled shells or those that contain 8 e- *electronegative elements gain e- to fill outer shell *electropositive elements donate e- to obtain filled outer shell Every bond is an exothermic reaction. We lose energy in bonds because energy + attraction are being concentrated and becoming denser, thus entropy decreases The possibilities for carbon are basically endless. Located in group 4, it has the ability to bond to almost anything and create any kind of structure. Atomic orbitals: how molecules get their shape S-orbitals: spherical p-orbitals: dumbbell shaped d-orbitals: Electron configuration rules: 1) Aufbau principle: orbitals are filled only after lower energy ones are filled 2) Pauli exclusion principle: max 2 electrons per orbital, with opposite spins 3) Hund’s rule: electrons of the same energy remain unpaired for as long as possible in parallel spins. Note: unpaired electrons try to get as far away as possible, which is why forming bonds are exothermic. VALENCE BOND THEORY Valence bond theory: a covalent bond forms when two atoms approach each other and a singly occupied orbital overlaps another one, thus pairing together and becoming attracted to both nuclei. The principle of maximum overlap: Orbitals with the greater directional character form stronger bonds because there is more electron density between atoms  The shape of the molecule is the result of attractive and repulsive forces acting on atoms that are either sp, sp^2, or sp^3 hybridized.  Bonds that look like cylinders head on are called sigma-bonds  Bonds that don’t are called pi-bonds Carbon can be tetrahedral (tetravalent) due to hybridization: Pauling’s mathematical model explaining how s and p orbitals can combine to distribute electron density. Steric effect: due to the size of the attached groups, there will also be repulsion. Electron pairs repel more than bonds. Pi-bonds: when double and triple bonds form. They bond sideways and are from overlapping p-orbitals. They are not as strong as sigma bonds because they are not as electronically dense, but do significantly strengthen the electron density between atoms. POLAR COVALENCE -joining atoms of different electronegativity Expressing polarity: covalent polar covalent ionic Lewis dot structure: Kekele structure: Dipole moment ???? = ???? × ???? *polar bonds may lead to polar Q = charge molecules. r = distance Dipole moment of molecules is the vector sum of bond dipole moment. These dipoles create the possibility of forming strong sigma-bonds Formal charge: the difference in the number of electrons in the valence shell as compared to the neutral atom. Isoelectronic: of having the same electronic configuration as another molecule or atom. Five Rules to Resonance: 1) Individual resonance forms are imaginary, not real. 2) Resonance forms differ only in the placement of their pi or nonbonding electrons. 3) Different resonance forms of a substance don’t have to be equivalent. 4) Resonance forms obey normal rules of valency. 5) The resonance hybrid is more stable than any individual resonance. There will be delocalization of electrons regardless of n, where n = 0,1, or 2 electrons. ACIDS AND BASES Lewis acid-base reactions involve dipoles to some extent.  Acid: species that accepts e- pair from a sigma bond  Base: species that donates e- pair from a sigma bond A neutral molecule has no formal charge We don’t have to have charged species to be a Lewis acid: hybridized sp^2, which a vacant p-orbital. Boron trihydride is a classic example of a Lewis acid. Ammonia (above) is hybridized sp^3, and is a classic Lewis base. As BH3 bonds, it must hybridize to sp^3 to get an additional sigma bond. Bronsted-Lowry  Acid: species that donates proton (H+)  Base: species that accepts proton (H+) Equilibrium always favors production of a weaker acid/base: ???????? ???? −???????????????? ???? Ex: ethanol (pKa=16) vs HCl (pKa=-7) The lower the pKa, the stronger the acid. In general: + − ???????? + ????:→ ???? ???? + ???? We usually determine the strength of an acid or a base aqueously, with respect to water. + − ???????? + ???? ????2↔ ???? ???? 3 ???? Thus ???????? = −log ???? ???? 3 + ]and ???? 3 +] = ???? − ] Stability of conjugate base is correlated to the position of equilibrium/strength of the acid as electron density plays a huge role. Three factors affecting stability: 1) Increasing electronegativity or electron density of acid increases the stability of the conjugate base 2) Increasing the size of the acid increases the stability of the acid, assuming relative electronegativity 3) Resonance can stability the conjugate base, which increases the strength of the acid. ALKANES AND CONFIGURATION Functional groups: a group of atoms that show characteristic chemical behavior and are relatively independent of their molecular environment. Three classes of functional groups 1) Carbon-carbon single, double, and triple bonds, and arenes 2) Polar sigma-bond: carbon bonded to more electropositive or electronegative element 3) Multiple bonds between carbon and a more electronegative element These groups are everywhere: take a biologically active compound, like a prostaglandin, for example. Alkanes are the least reactive of all functional groups, mainly because they contain a nonpolar sigma bond. They undergo mainly two different kinds of reactions: 1) Combustion (oxidation reaction) ∆ ???????? 3???? ????2 + 3???? → 3????2 + 4???? ????2+ ℎ????????????2 2) Halogenation (free radical halogenation) – a substitution reaction ???????? ???????? + ???????? → ℎ???? ???????? ???????? ???????? ???????? + ???????????? 3 3 2 3 2 General formula for alkanes: ???????????? ????????+???? Many different structures are possible for alkanes, which have the same molecular formula. For instance: C6H14. These are isomers: same molecular formula, different structural properties Stereochemistry: the study of molecules in three dimensions  Configurational isomers: isomers having different bond connectivity  Stereoisomers: isomers having the same bond connectivity, but different projection in space o Configurational/Constitutional stereoisomers: Interconversion involves breaking sigma- bonds (see above) o Geometric stereoisomers: Breaking pi-bonds o Conformational stereoisomers: can be rotated around sigma-bonds Need to know: how to draw a Newman diagram. Staggered: when sigma bonds are not directly in line with one another. (60 degrees between them) Eclipsed: when sigma bonds are directly in line with one another. (0 degrees between them) Values to memorize: Strain kJ H-H torsional strain 4 H-methyl torsional strain 6 Methyl-methyl torsional + steric strain 11 Methyl-methyl steric strain 3.8 Kinds of strain and where they come from: Strain How Torsional (kJ) electron density repulsion Steric (kJ) repulsion due to size of group Angle (degrees, excess energy as a result in kJ) angle bent out of ideal, Total Calculated from combustion reaction Some special names for n-butane when doing conformational analysis down the 2-3 carbon bond:  Anti-conformation: when the CH3s are as far away as possible. 0 kJ total strain.  Gauche formation: when the CH3s are staggered but not 180 degrees away from each other, creating major torsional and steric strain, which is 19 kJ in total. CYCLOALKANES AND CONFIGURATION Here we start to consider angle strain as well as torsional and steric strain. However, we only have to be cognizant of the fact that the energy cost of angle strain is caused by the fact that there is not enough overlapping electron density, and thus it takes more energy to form the bond. Puckering: shifting bond angle out of the perfectly planar angle and into one where there may be greater angle strain but reduced torsional strain for an over reduced total strain.  Cyclobutane  Cyclopentane  Cyclohexane


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