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by: Alice Hsu

CHEM 51 WEEK 1 NOTES Chemistry 51

Marketplace > Dartmouth College > Chemistry > Chemistry 51 > CHEM 51 WEEK 1 NOTES
Alice Hsu
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Chemistry 51 week 1 notes. Organic chemistry week 1, detailing a brief review of general chemistry. We cover structural theory, a simplistic model of an atom, atomic bonding, valence theory, orbita...
Organic Chemistry
Peter Jacobi
Class Notes
Organic Chemistry, Organic Chem, General Chemistry, Chemistry, acids and bases, Resonance Forms, atomic orbitals, bond polarity and atom formal charges




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Popular in Chemistry

This 6 page Class Notes was uploaded by Alice Hsu on Sunday September 18, 2016. The Class Notes belongs to Chemistry 51 at Dartmouth College taught by Peter Jacobi in Fall 2016. Since its upload, it has received 7 views. For similar materials see Organic Chemistry in Chemistry at Dartmouth College.


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Date Created: 09/18/16
CHEM 51 WEEK 1 INTRO TO ORGO The beginning of organic chemistry began with the formation of earth’s primordial atmosphere, when CO, NH3, CO2, HCN collided under a lot of pressure, heat, and energy, forming purines and pyrimidines, the basis of life, and amino acids, which formed peptides, proteins, and enzymes. 1828: Fredrich Wohler proved that the theory behind “vital force” was incorrect and that one could synthesize organic compounds from inorganic compounds: This lead to a structural theory—how atoms are put together to form molecules. 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. e.g. methane is a perfect tetrahedron: carbon is sp^3 hybridized and forms sp^3-s sigma bonds with hydrogen Ethane: the combination of methyl radicals, which are both sp^3. the sp^3-sp^3 sigma-bond is weaker because the electron density is less concentrated than an sp^3-s sigma-bond Steric effect: due to the size of the attached groups, there will also be repulsion. Electron pairs repel more than bonds: EX: NH3 Is also sp^3 hybridized, but has a lone pair of electrons, making extra repulsion and minimizing the N-H bond angle. N-H is also a shorter bond compared to C-H in methane because N is more e-neg Oxygen, which has two lone pairs, continues the trend, with an even smaller O-H bond length and angle. Pi-bonds: when double and triple bonds form. They bond sideways and are from overlapping p-orbitals. EX: ethylene, which is sp^2 hybridized, meaning three sp^2 orbitals, and one 2p_z orbital These pi-bonds make the C-C bond shorter and stronger by increasing the overall electron density, but pi-bonds are not as strong as sigma-bonds because they have nodes. EX: acetylene continues the trend with a triple bond acetylene is sp-hybridized. The electron density around the C-C bond is like a cylinder, or a donut coming out the page. This C-C triple bond is the strongest in O chem. POLAR COVALENCE -joining atoms of different electronegativity Expressing polarity: covalent polar covalent ionic Lewis dot structure: Kekele structure: Dipole moment ???? = ???? × ???? Q = charge r = distance *polar bonds may lead to polar molecules. Dipole moment of molecules is the vector sum of bond dipole moment. Ex: chloromethane Ex: ammonia, the lone e- pair creates a net dipole Ex: water Ex: dichloromethane, a net dipole moment Ex: lithium methyl Ex: carbotetrachloride, no dipole because of the symmetry 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. 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. Sp^3-s Sp^3-Sp^3 sigma bonds sigma bonds Summing up resonance from last time: There will be delocalization regardless of n, where n = 0, 1, or 2 electrons 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.


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