CHE-300 Organic Chemistry I Dr. Noel Sturm; office: NSM D-323

CHE-300 Organic Chemistry I Dr. Noel Sturm; office: NSM D-323
Dr. Noel Sturm; office: NSM D-323 (310) or office hours: by Appointment

Print the slides as Hand-Outs, 6 per page, before coming to class.
Web page: Print the slides as Hand-Outs, 6 per page, before coming to class.

Text: Organic Chemistry, Morrison & Boyd (6th Edition) Study Guide to..., Morrison & Boyd Supplement... Model kit

Grading: No Curve!, Everyone Can Earn An “A”
A A B B B C C C D D F 59-0 Weekly Exams No make ups! Begin at 10:00!

Daily Homework: Required!
Final Exam, June 27th, 10-12:15, NSM C-221 comprehensive!! Cheating: Don’t do it! The penalties are severe. Turn off all cell phones and pagers!

Organic Chemistry; difficult, challenging. “memorization course” (NOT
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. participate in “quick” thinks 4. complete the work-sheets 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 = 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. % H H C C C Cl Cl F

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:
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 => 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 sp o 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. O :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

CO :O=C=O: sp linear vector sum = 0 non-polar molecule H2O .. H—O—H AX2E 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 polar covalent => dipole-dipole & Hydrogen bonding ionic bonding => ionic attractions

bonding => shape => polarity => 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  CH Mg(NH2)Br NaOH H2SO4  H2O NaHSO4 base acid acid base base acid acid base

- + BF3 + :NH3  F3B:NH3 Lewis Lowry-Brønsted
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. - + BF :NH3  F3B:NH3 Lewis Lowry-Brønsted

Rule: acid/base reactions must run “down hill.”
stronger acid/base  weaker acid/base H2SO H2O  HSO H3O+ stronger stronger weaker weaker acid base base acid H2O NH3  NH 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

├────┼────┐ ┌────┬────┬────┬────┬────┼────┤

Which is the stronger acid?
H2O or H2S? What is the order of base strength? F Cl Br I- oxygen & sulfur are in the same family and sulfur is bigger: H2S > H2O in the halogen family base strength decreases with increasing size: F- > Cl- > Br- > I-

Will H2O react with NaSH as shown below?
H2O NaSH  NaOH H2S Will the following reaction proceed as shown? HI NaCl  HCl NaI WA SA no, H2O < H2S SA WA yes, HI > HCl

Isomers - different compounds with the same molecular formula.
example: C2H6O CH3CH2OH CH3OCH3 ethyl alcohol dimethyl ether bp 78oC bp –24oC

CHE-300 Organic Chemistry I Dr. Noel Sturm; office: NSM D-323

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