CHE-300 & 302 Organic Chemistry I & II Dr. James Lyle; office: NSM D-323

CHE-300 & 302 Organic Chemistry I & II Dr. James Lyle; office: NSM D-323 (310) 243-3388 or 243-3376 jlyle@csudh.edu office hours: M-Th, 8:00 am – 10:00 am

Web page: http://chemistry.csudh.edu/

texts: Organic Chemistry, Morrison & Boyd (6th) Supplement to..., Morrison & Boyd (optional) Supplement... Model kit

Grading: traditional, no curve! A=100%-93%, A-=92%-90%,B+=89%-88%, B=87%-83%,etc. Daily exams = 75% Final exam = 25% *** Daily exams No make ups! Drop two lowest scores. Begin at 10:00!

Daily Homework:Required! (hold until called for) Final Exam, July 7, 10-12:00 comprehensive, difficult! Cheating: Don’t do it! The penalties are severe. Turn off all cell phones and pagers!

Organic Chemistry; difficult, challenging! “memorization course” (NOT! well…maybe), body of knowledge + application of theory! How to succeed? 1. look over the text before lecture. 2. listen carefully to lectures 3. read the text (take notes) 4. do the homework (twice...?) 5. review

Organic Chemistry - the study of the compounds of carbon, their properties and the changes that they undergo. Descriptive approach - nomenclature syntheses reactions mechanisms ...

First: review topics from gen. chem. important to o-chem. atomic structure subatomic particles: mass charge protons neutrons electrons nucleus: protons & neutrons electron shells & subshells: electrons 1 amu +1 1 amu 0 ~0 amu - 1

atomic number = number of protons in the nucleus of the atom (different for each element); Hydrogen = 1, Helium = 2, Lithium = 3,... [also the number of electrons in a neutral atom] Iron = 26 26 protons = +26 26 electrons=-26 net charge= 0

atomic mass = mass of an atom; sum of the weights of the protons & neutrons. But, not all atoms of a given element are identical. isotopes- atoms of the same element with different numbers of neutrons.

examples of isotopes prot. neut. % H1 1 0 99.985 H2 1 1 0.015 C12 6 6 98.89 C13 6 7 1.11 C14 6 8 ... Cl35 17 18 75.53 Cl37 17 20 24.47 F19 9 10 100

atomic weight: weighted average mass of the atoms; combining weight... electrons => energy shells & subshells about the nucleus. shells = 1, 2, 3, 4, ... subshells = s, p, d, f orbitals = region in space where an electron of given energy is likely to be found; no more than two electrons of opposite spin per orbital (Pauli exclusion principle).

maximum number of electrons per subshell: s 2 p 6 d 10 f 14

order of filling 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f spectral notation: 1s2,2s2,2p6…

Fluorine (at.# 9) 9p/9e 1s2,2s2,2p5 Chlorine (at.# 17) 17p/17e 1s2,2s2,2p6,3s2,3p5 Bromine (at.# 35) 35p/35e 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p5 Iodine (at.# 53) 53p/53e 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p5

valence electrons = electrons in the outermost shell Fluorine has 7 valence elect. Chlorine has 7 valence elect. Bromine has 7 valence elect. Iodine has 7 valence elect.

PERIODIC CHART OF THE ELEMENTS I VIII ┌────┐ ┌────┐ │ H │ │ He │ │ 1 │ II III IV V VI VII │ 2 │ ├────┼────┐ ┌────┬────┬────┬────┬────┼────┤ │ Li │ Be │ │ B │ C │ N │ O │ F │ Ne │ │ 3 │ 4 │ │ 5 │ 6 │ 7 │ 8 │ 9 │ 10 │ ├────┼────┤ ├────┼────┼────┼────┼────┼────┤ │ Na │ Mg │ │ Al │ Si │ P │ S │ Cl │ Ar │ │ 11 │ 12 │ │ 13 │ 14 │ 15 │ 16 │ 17 │ 18 │ ├────┼────┼────┬────┬────┬────┬────┬────┬────┬────┬────┬────┼────┼────┼────┼────┼────┼────┤ │ K │ Ca │ Sc │ Ti │ V │ Cr │ Mn │ Fe │ Co │ Ni │ Cu │ Zn │ Ga │ Ge │ As │ Se │ Br │ Kr │ │ 19 │ 20 │ 21 │ 22 │ 23 │ 24 │ 25 │ 26 │ 27 │ 28 │ 29 │ 30 │ 31 │ 32 │ 33 │ 34 │ 35 │ 36 │ ├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤ │ Rb │ Sr │ Y │ Zr │ Nb │ Mo │ Tc │ Ru │ Rh │ Pd │ Ag │ Cd │ In │ Sn │ Sb │ Te │ I │ Xe │ │ 37 │ 38 │ 39 │ 40 │ 41 │ 42 │ 43 │ 44 │ 45 │ 46 │ 47 │ 48 │ 49 │ 50 │ 51 │ 52 │ 53 │ 54 │ ├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤ │ Cs │ Ba │ La │ Hf │ Ta │ W │ Re │ Os │ Ir │ Pt │ Au │ Hg │ Tl │ Pb │ Bi │ Po │ At │ Rn │ │ 55 │ 56 │ 57 │ 72 │ 73 │ 74 │ 75 │ 76 │ 77 │ 78 │ 79 │ 80 │ 81 │ 82 │ 83 │ 84 │ 85 │ 86 │ ├────┼────┼────┼────┼────┼────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘ │ Fr │ Ra │ Ac │ │ │ │ 87 │ 88 │ 89 │104 │105 │ └────┴────┴────┴────┴────┘ ┌────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┬────┐ │ Ce │ Pr │ Nd │ Pm │ Sm │ Eu │ Gd │ Tb │ Dy │ Ho │ Er │ Tm │ Yb │ Lu │ │ 58 │ 59 │ 60 │ 61 │ 62 │ 63 │ 64 │ 65 │ 66 │ 67 │ 68 │ 69 │ 70 │ 71 │ ├────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┼────┤ │ Th │ Pa │ U │ Np │ Pu │ Am │ Cm │ Bk │ Cf │ Es │ Fm │ Md │ No │ Lr │ │ 90 │ 91 │ 92 │ 93 │ 94 │ 95 │ 96 │ 97 │ 98 │ 99 │100 │101 │102 │103 │ └────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┴────┘

periodic chart of the elements metals & nonmetals families (groups) of elements alkali metals (group I) Li,Na,K,... alkaline earths (group II) Be,Mg,Ca,... halogens (group VII) F,Cl,Br,I,... noble gases (group VIII or 0) He,Ne,Ar,... group number = valence elec.

Chemical bonding (classical) chemical bond: force that holds atoms together in compounds. ionic bond ~ between metals & non-metals covalent bond ~ between non-metals & non-metals

definitions: ionic bond: a chemical bond formed by the transfer of valence electrons to achieve noble gas electron config-urations, resulting in ions held together by electrostatic attraction. covalent bond: chemical bond formed by the sharing of valence electrons to achieve noble gas electron configurations.

ionic bond example: sodium chloride sodium = Na, atomic # 11 1s2,2s2,2p6,3s1 neon = Ne, atomic # 10 1s2,2s2,2p6 if Na loses 1 elect. then it will have a noble gas elect. config. like Ne but will be charged, +1 ( 11p/10e ). => Na+ sodium ion

chlorine = Cl, atomic # 17 1s2,2s2,2p6,3s2,3p5 argon = Ar, atomic # 18 1s2,2s2,2p6,3s2,3p6 if chlorine can gain an electron it will have a noble gas electron config. like argon but will be charged -1 (17p/18e) Cl- sodium chloride = NaCl or Na+Cl-

covalent bonds Lewis Dot representations H Be :Cl Ne C O H2O = H:O:H see homework! review your gen chem text! .. . . . . .. .. . . . . . .. . .. ..

.. .. .. .. CO2 :O::C::O: :O=C=O: N2 :N:::N: :NN: HCN H:C:::N: H-CN: .. .. H2CO H:C::O: H-C=O: .. | H H

atomic orbitals s p d etc.

hybrid atomic orbitals s + p => 2 sp hybrids s + p + p => 3 sp2 s + p + p + p => 4 sp3

Hybrid atomic orbitals: sp = linear; 180o sp2 = trigonal; 120o sp3 = tetrahedral; 109.5o

VSEPR (valence shell electron pair repulsion) prediction of hybridization number of ligands (X) plus number of unshared pair of valence electrons (E) equals number of orbitals needed  what type of hybrid orbitals are needed

.. .. eg. H2O => H:O:H or H—O—H 2 ligands + 2 lone pair = 4 orbitals AX2E2 sp3 tetrahedral, 109.5o water is a bent molecule with bond angles of 105o .. ..

VSEPR AX2 sp 180o linear AX3 sp2 120o trigonal AX2E sp2 ~120o or “bent” AX4 sp3 109.5o tetrahedral AX3E sp3 ~109.5o or “pyramidal AX2E2 sp3 ~109.5o or “bent”

We can use the VSEPR method to predict the shape and bond angles for simple covalent molecules. SHAPE is important! review gen chem text! Do the homework!!!!!

Polarity Covalent bonds are polar when the two atoms sharing electrons have different electronegativities. eg. H—Cl δ+ δ- a charge separation or a dipole gives a polar bond. .. .. O2 :O=O: has a non-polar bond

Representation of dipoles using vectors a) magnitude = length b) direction = positive  negative A molecule will be non-polar if the vector sum of the bond dipoles is zero; eg. they cancel one another. A molecule with be polar if the vector sum of the bond dipoles is non-zero.

Determining polarity of covalent molecules: • Lewis dot structure • VSEPR  hybridization  shape of the molecule • dipoles for polar bonds • vector sum of the bond dipoles • vector sum = 0  non-polar molecule • vector sum  0  polar molecule

CO2 :O=C=O: sp linear vector sum = 0 non-polar molecule H2O .. H—O—H AX2E2 sp3 tetrahedral (bent) .. vector sum  0 polar molecule!

CH3OH Both C & O are sp3 hybridized. The bond dipole vectors do not cancel each other and the molecule is polar. NB: must know shape to determine polarity!

Intermolecular forces. Attractions between molecules. ionic attractions Na+Cl- (very strong) Cl-Na+ dipole-dipole attractions H—Br Br—H hydrogen bonding ( H attached to N,O,F ) H—O----H—O | | H H van der Waals (London forces) Br—Br (weak) Br—Br

intermolecular attractions strongest ionic attractions dipole-dipole / hydrogen bonding van der Waals weakest ionic bonds => ionic attractions polar covalent => dipole-dipole attractions non-polar covalent => van der Waals

Cl2 CO2 H2O CH4 KBr non-polar covalent => van der Waals non-polar covalent => van der Waals polar covalent => dipole-dipole & Hydrogen bonding non-polar covalent => van der Waals ionic bonding => ionic attractions

bonding => shape => polarity => physical properties physical properties: melting point boiling point solubility The stronger the intermolecular forces the higher the mp/bp. Ionic substances have significantly higher mp/bp than do covalent substances. [note: mp/bp also increase with increasing size.]

Prediction of mp/bp (relatively high or low?): Mg(OH)2 CH3OH CH2O CH3CH3 mp bp 350oC -- -94oC 65oC -920C -21oC -183oC –89oC ionic => ionic attractions polar => dipole-dipole + H-bond polar => dipole-dipole non-polar => van der Waals

Solubility “like dissolves like” ~ water soluble? must be ionic or highly polar + H-bond (hydrophilic) ~ water insoluble? must be non-polar or weakly polar (hydrophobic) Most organic compounds are water insoluble!

Acids/Bases historic: acids – from L. acidus = “sour” sour taste react with metals  H2 react with bases  water + salts change litmus  red react with limestone  CO2 examples: HCl, H2SO4, HNO3, HClO4

historic: bases - bitter taste soapy feel react with acids  water + salts change litmus  blue examples: NaOH, Al(OH)3, K2CO3, NaHCO3

Lowry-Brønsted Acid - a substance that donates a proton (H+) in a chemical reaction. Lowry-Brønsted Base – a substance that accepts a proton (H+) in a chemical reaction. CH3MgBr + NH3 CH4 + Mg(NH2)Br NaOH + H2SO4  H2O + NaHSO4 base acid acid base base acid acid base

Lewis Acid – a substance that accepts an electron pair in a chemical reaction to form a covalent bond. Lewis Base – a substance that donates an electron pair in a chemical reaction to form a covalent bond. - + BF3 + :NH3 F3B:NH3 Lewis Lowry-Brønsted

Rule: acid/base reactions must run “down hill.” stronger acid/base  weaker acid/base H2SO4 + H2O  HSO4- + H3O+ stronger stronger weaker weaker acid base base acid H2O + NH3  NH4+ + OH- weaker weaker stronger stronger acid base acid base (note direction of reactions)

Within a period of the periodic chart, acid strength increases with increasing electronegativity: CH4 < NH3 < H2O < HF Within a family of elements, acid strength increases with increasing size: HF < HCl < HBr < HI

CHE-300 & 302 Organic Chemistry I & II Dr. James Lyle; office: NSM D-323