CHEMISTRY 31 Organic Chemistry Lecture I Dr. John D. Spence [email protected] jspence@csus. edu http://www.csus.edu/indiv/s/spencej What is Organic Chemistry? 1780’s Chemistry of compounds from living organisms Vitalism: 1750-1850 Inorganic compounds – non-living source Organic compounds – living source August Kekule (1861) Organic chemistry = a study of the compounds of carbon Origins of Life Stanley Miller 1950 Syyyppynthesis of amino acids, formaldehyde, purines and pyrimidines from earth’s primordial atmosphere (H2, H2O, O2, NH3, CO2, CH4) and static discharge 1 Organic Chemist’s View of 6 H C 1.007 the Periodic Table of 12.01 94 Elements 1 3 6 5 6 7 8 9 6 Li C B C N O F C 6.941 12.01 10.81 12.01 14.00 15.99 18.99 12.01 1 1 1 67 94 840 1 11 12 13 14 15 16 17 6 N M Al Si P S Cl C 26.98 28.08 30.97 32.06 35.45 12.01 a g 154 55 376 6 27 1 22.98 24.30 977 5 19 20 6 22 6 24 25 26 6 6 29 30 6 6 6 6 35 6 K Ca C Ti C Cr Mn Fe C C Cu Zn C C C C Br C 39.09 40.07 12.01 47.88 12.01 51.996 54.938 55.847 12.01 12.01 63.546 65.39 12.01 12.01 12.01 12.01 79.90 12.01 83 8 1 1 1 0 1 1 1 1 1 1 4 1 6 6 6 6 6 6 6 6 6 46 6 6 6 6 6 6 53 6 C C C C C C C C C Pd C C C C C C I C 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 106.42 12.01 12.01 12.01 12.01 12.01 12.01 126.9 12.01 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 045 1 6 6 6 6 6 6 6 76 6 6 6 80 6 6 6 6 6 6 C C C C C C C Os C C C Hg C C C C C C 12.01 12.01 12.01 12.01 12.01 12.01 12.01 190.2 12.01 12.01 12.01 200.59 12.01 12.01 12.01 12.01 12.01 12.01 1111111 111 111111 6 6 6 6 6 6 6 6 6 110 111 112 114 116 118 C C C C C C C C C (269) (272) (277) (285) (289) (293) 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 1 1 1 1 1 1 1 1 1 *Lanthanide Series 6 6 6 6 6 6 6 6 6 6 6 6 6 6 C C C C C C C C C C C C C C 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 1 1 1 1 1 1 1 1 1 1 1 1 1 1 †Actinide Series 6 6 6 6 6 6 6 6 6 6 6 6 6 6 C C C C C C C C C C C C C C 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 12.01 1 1 1 1 1 1 1 1 1 1 1 1 1 1 Atomic Structure Structure of an atom Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m) Negatively charged electrons are in a cloud (10-10 m) around nucleus Diameter is about 2 10-10 m (200 picometers (pm)) [the unit angstrom (Å) is 10-10 m = 100 pm] 6 C Atomic Number and Atomic Mass 12. 011 The atomic number (Z))p is the number of protons in the atom's nucleus (neutral atom then has same number of electrons) The mass number (A) is the number of protons plus neutrons All the atoms of a given element have the same atomic number Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numb(bers (13C, 2HD)H or D) The atomic mass (atomic weight) of an element is the weiggghted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes Atomic Structure: Orbitals Quantum mechanics: describes electron energies and locations by a wave equation Wave function solution of wave equation Each wave function is an orbital, A plot of 2 describes where electron most likely to be Electron cloud has no sppyecific boundary so we show most probable area Shapes of Atomic Orbitals for Electrons Four different kinds of orbitals for electrons based on those der ive d for a hy drogen a tom Denoted s, p, d, and f s and p orbitals most imppgortant in organic chemistr y s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle Orbitals and Shells Orbitals are grouped in shells of increasing size and energy Differen t s he lls con ta in differen t num bers an d kin ds o f or bita ls Each orbital can be occupied by two electrons First shell contains one s orbital, denoted 1s, holds only two electrons Second shell contains one s orbital (2s) and three p orbitals (2p), eig ht electrons Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons valence shell octet for second row p-Orbitals In each shell there are three ppperpendicular p orbitals, px, py, and pz, of equal energy Lobes of a p orbital are separated by region of zero electron density, a node Atomic Structure: Electron Configurat ions Ground-state electron configuration of an atom lists orbitals occupied by its electrons. Rules: 1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d (AfbAufbau (“bu ild-up”)i”) princ il)iple) 2. Electron spin can have only two orientations, up and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If t wo or more empt y orbit al s of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule). Electron Configuration and Bonding 2p 2s 1s Lithium Carbon Oxygen Fluorine Z = 3 Z = 6 Z = 8 Z = 9 How do atoms fill octet? BONDING Ionic = bonds from exchanging electrons Covalent = bonds from sharing electrons Va lency = num ber o f cova lent bon ds nee de d to fill val ence octet C = tetravalent N = trivalent O = divalent H = monovalent Drawing Structures and Bonds Lewis structures (()electron dot) show valence electrons of an atom as dots Hydrogen has one dot, representing its 1s electron Carbon has four dots (2s2 2p2) Kekule structures (line-bond structures) have a line drawn between two atoms indicating a 2 electron covaltbdlent bond. Stable molecule results at completed shell, octet ((geight dots) for main-ggproup atoms (two for h yg)ydrogen) (third row atoms can expand beyond octet) Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons Nitrogen atom in ammonia (NH3) Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair Bonding Model #1: Valence Bond Theory Singly occupied orbital on one atom overlaps a silingly occup idbitlied orbital on the other atom Electrons are paired and localized in the over lapp ing or bita ls an d are attracted to nuclei of both atoms H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals Certain optimal distance Smaller, more directional orbitals make stronger bonds due to better orbital overlap pingpong ball - pingpong ball pingpong ball - basketball pingpong ball - watermelon Bond Energy Reaction 2 H· H2 releases 436 kJ/mol Product has 436 kJ/mol less energy than two atoms: H–H has bond strength of 436 kJ/mol. (1 kJ = 0. 2390 kca l; 1 kca l = 4. 184 kJ) Bond Length Distance between nuclei that leads to maximum stability If too close, they repel because both are positively charged If too far apart, bonding is weak Hybridization for the Carbon Atom 2p 2s Ground state Four new sp3 orbitals Carbon Why Hybridize? 1. can now form 4 bonds instead of two to fill octet 2. more directional orbital (increased s-character) gives stronger bond from better overlap of orbital 3. maximize separation of electrons (VSEPR) sp3 Orbitals sp3 hybrid orbitals: s orbital and three p orbitals combine to form four equivalent, unsymmetrical, tetrahedral orbitals (sppp = sp3), Pauling (1931) The Structure of Methane sp3 orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds Each C–H bond has a strength of 438 kJ/mol and length of 109 pm Bond angle: each H–C–H is 109.5°, the tetrahedral angle. Hybridization: sp3 Orbitals and the Structure ofhf Ethane Two C’s bond to each other by overlap of an sp3 orbita l from eac h Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds C–H bond strength in ethane 423 kJ/mol C–C bond is 154 pm long and strength is 376 kJ/mol All bond angles of ethane are tetrahedral one Csp3-Csp3 bond 6C6 Csp3-H1s bdbonds Hybridization: sp2 Orbitals and the Structure ofhlf Ethylene sp2 hybrid orbitals: 2s orbital combines with two 2p orbitals, giving 3 orbitals (spp = sp2).