교과목명 유기화학(1) 학점 3 강의실 과B133 강의시간 수3,4 ......연습문제및출석및수업참여도: 50점 성적평가방법 강좌운영방식 교재중심강의

Organic Chemistry John McMurry 6th Ed교재 및 참고문헌

중간고사(2회): 100점 기말고사(1회): 100점연습문제 및 출석및 수업참여도: 50점

성적평가방법

교재중심강의강좌운영방식

일반화학1,2선수과목(선수학습)

탄소화합물의 구조, 명명법, 그리고 기초반응을 습득한다.

수업목표 및 개요

의예과 2학년수강대상

[email protected]면담시간

2123-2644연락처과학관410호연구실

이과대학 화학소속전철호교수명

수3,4, 금4강의시간과B133강의실

3학점유기화학(1)교과목명

학기말고사6.13 - 6.16 16

Exam III 6.06 - 6.11 15

Ch.12 Mass and Infrared Spectroscopy5.30 - 6.04 14

Ch. 11.11 The E2 Elimination5.23 - 5.28 13

Ch.11 Reaction of Alkyl Halides5.16 - 5.21 12

Ch.10 Alkyl Halides5.9 - 5.14 11

Exam IICh.9 Stereochemistry5.2 - 5.7 10

Ch.8 Alkynes: Introduction to Organic Synthesis4.25 - 4.30 9

중간고사4.18 - 4.23 8

Ch.7 Alkenes: Reactions and Synthesis4.11 - 4.16 7

Ch.6 Alkenes: Structure and Reactivity4.04 - 4.09 6

Exam ICh.5 An Overview of Organic Reactions3.28 - 4.02 5

Ch.4 Stereochemistry of Alkanes and Cycloalkanes3.21 – 3.26 4

Ch.3 Organic Compounds: Alkanes and Cycloalkanes3.14 – 3.19 3

Ch.2 Polar Bonds and Their Consequences3.7 – 3.12 2

Ch.1 Structure and Bonding3.2 – 3.5 1

Exam수업내용기간주

Ch.1 Structure and Bonding

mid-1700s: alchemist, no differences between substances obtained from living sources and those obtained from minerals

Organic Chemistry?

1770, Torbern Bergman: first express the difference of organic/inorganic; organic chemistry = chemistry of compounds found in living organisms

organic ~ vital force

1816, Michael Chevreul: for the first time, one organic substance was converted into others

Animal fatNaOH

H2OSoap + Glycerin

Soap H3O+"Fatty acids"


1828, Friedrich WŐhler: convert "inorganic" salt to "organic" substance

NH4+ -OCN

heatH2N

CNH2

O

Ammonium cyanate Urea

mid-1800s: no more vitalistic theory

~ today: no clear distinction between organic and inorganic chemistry; unified chemistry; but, all organic chemicals contain the element carbon


- the chemistry of compounds found in living organisms- the study of carbon compounds

H C N O FPSi S Cl

BrI

Organic Chemistry?

- DNA, RNA, amino acid, proteins, carbohydrates……- medicine, dyes, polymers, plastics, food additives, pesticides, fragrant…..

1.1 Atomic Structures

nucleus (protons + neutrons): dense, positively charged; occupies most of atomic mass

orbiting electrons: negligible mass, negatively charged

d= 10-14~10-15m

~10-10m (1 A)


atomic number (Z): # of protonsmass number (A): # of protons and neutrons (neutrons are neutral)

for neutral atomsatomic number = # of protons (positively charged) = # of electrons (negatively charged)

1H: 1 (proton) + 0 (neutron) = 12He: 2 (proton) + 2 (neutron) = 4

1.008 for hydrogen: 1H (99.985%, 1.0078), 2H (0.015%, 2.0141)12.011 for carbon: 12C (98.89%, 12.000), 13C (1.11%, 13.0034)


Isotope: same atomic number but different mass number (different number of neutrons)

All the atoms in a given element have the same atomic number but they can have different mass numbers depending on how many neutrons they have.

Atomic weight: the weighted average atomic mass units of an element’s isotopes

- electron distribution: quantum mechanics- wave equation solution: wave function (ψ) or orbital: wave function of one electron atom

s orbital p orbital d orbitals

1.2 Atomic Structure: Orbitals


shapes of orbitals: s, p, d, f

ψ2 ; the volume of space around a nucleus where an electron can most likely be found (probability)

shells: atomic orbitals with the same principal quantum number

1s

2s

2p

3s

3p

3d

ener

gy

1st shell

2nd shell

3rd shell(18 e-)

(8e-)

(2 e-)


- distribution of electrons in an atom: each orbital can occupy 2 e-

shapes of 2 p orbitals

2px orbital

x

y

z

2py orbital

x

y

z

2pz orbital

x

y

z


node: a region of zero electron densitynodal plane:

1.3 Atomic Structure: Electron Configuration


ground-state electron configuration: lowest-energy arrangement

rule 1: The lowest-energy orbitals fill up first (Aufbau principle)

1s 2s 2p 3s 3p 4s 3d

1s

2s

2p

Energy


rule 2: Only two electrons can occupy an orbital and they must be of opposite spin (Pauli exclusion principle): electrons spins (↑) or (↓)

1s

2s

2p


If two or more orbitals of equal energy are available, one electron occupies each until all orbitals are half-full. Only then does a second electron occupy one of the orbitals (Hunt’s rule): The electrons in the half-filled orbitals all have the same spin

rule 3:

1s

2s

2p

1sHydrogen

1atomic number

1s

2s

2pCarbon

6atomic number


ground-state electron configurations of some elements

1s

2s

2pOxygen

8atomic number


ground-state electron configuration of oxygen


Practice

Chloride: atomic number 17

1s

2s

2p

3s

3p

1.4 Development of Chemical Bonding Theory

In 1858, Kekule-Couper: tetravalent carbon

In 1865, Kekule-Couper: multiple bonding, ring systems

In 1874, van't Hoff: tetrahedral carbon (3-dimensional)

(2-dimensional)

Cin plane

dashed line (behind the plane)

heavy wedged line (out of the plane)


1.5 Covalent Bonds

Why do atoms bond together?


How can bonds be described electronically?difficult to answer

Because the compound is more stable, has less energy than the separate atoms

Octet rule: 8 electrons in valence shell, special stability; noble-gas configuration

Ne (2 + 8); Ar (2 + 8 + 8); Kr (2 + 8 + 8 + 8)

Na Na+ + e-

1s2 2s2 2p6 3s1 1s2 2s2 2p6


group 1A (alkali metal): loose 1 electron

group 7A (halogen): gain 1 electron

Cl Cl-+ e-

1s2 2s2 2p5 1s2 2s2 2p6


Cl-Na+

ionic bond: held together by electrostatic attraction

C C4+ + 4 e-

1s2 2s2 2p2 1s2X

How about the middle elements?

- satisfy octet by sharing electrons: covalent bond


1s1 2s2 2p2

H C

2s2 2p3

N

2s2 2p4

O

2s2 2p5

F

molecule: the neutral collection of atoms held together by covalent bonds

Lewis structure or electron-dot structure:

- valence electrons are represented by dots


C

H

H H

H

OH

H

N

H

H H FH

CH4 H2O NH3 HF

C

H

H

H

O

H

CH3OH


H Cl

Br FO

N

B

C

one bond two bond

three bond four bond


nonbonding electrons (lone-pair electrons); valence electrons that are not used for bonding

N

H

H HNH3 H N HH

Kekule structure (line-bond structure)


CH4 H2O NH3 CH3OH

CH

H HH

OH H NH

H H CH

H OH

H

C

H

H H

H

OH

H

N

H

H H FH

CH4 H2O NH3 HF

Lewis structure

Kekule' structure


PH?

Practice

GeCl?

AlCl? CH?Cl?

PH3PH

HH

P : 5A

GeCl4GeCl

ClCl

Cl

Ge : 4A

AlCl3AlCl

ClCl

Al : 3A

CH2Cl2CH

HCl

Cl

C : 4A

1.6 Valence Bond Theory and Molecular Orbital Theory

How does the electron sharing occur?


two models for covalent bond formation

valence bond theory

"electrons are localized around the bond"

molecular orbital theory

"electrons are delocalized over the molecule"

explains wellσ-bonds

explains wellπ-bonds


Valence Bond Theory

key ideas

• Covalent bonds are formed by overlap of two atomic orbitals, each of which contains one electron. The spins of the two electrons are opposite.

• Each of the bonded atoms retains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.

• The greater the amount of orbital overlap, the stronger the bond.

H H+ H H

1s H2 molecule1s


H H

circular cross-section

- H-H bond is cylindrically symmetrical

sigma (σ) bond: bonds formed by the head-on overlap of two orbitalsalong a line drawn between the nuclei

too close: nuclei repeltoo far: unable to share electrons→ optimum distance for maximum stability: bond length


ener

gy

H2 molecule

two hydrogen atoms

436 kJ/mol (104 kcal/mol)

2 H H2

released when bond formedabsorbed when bond breaks

How close are the two nuclei in the H2 molecule?


ener

gy 0

+

_

HH (too close)

H H

H H (too far)

74 pmbond length

A plot of energy versus internuclear distance for two hydrogen atoms.


1.12 Molecular Orbital Theory

key ideas

• Molecular orbitals describe regions of space in a molecule where electrons are most likely to be found, and they have a specific size, shape, and energy level.

• Molecular orbitals are formed by combining atomic orbitals. The number of MO's formed is the same as the number of atomic orbitalscombined.

• bonding MO: lower in energy than the starting atomic orbitalsantibonding MO: higher in energynonbonding MO: same energy


ener

gy

H 1s orbital H 1s orbital

H H antibonding MO (unfilled)

H H bonding MO (filled)

H H

H H

1.7 Hybridization: sp3 Orbitals and the Structure of Methane

Hydrogen molecule is simple, but complicate for tetravalent carbon atom


Carbon uses two kinds of orbitals (2s22p2) to form bondsbut four identical bonds (tetrahedral). How?

Hybridization (1931, Linus Pauling): mathematically explained how an s orbital and three p orbitals can combine (sp3 orbitals)

- valence bond theory

2s 2py

2pz2px

hybridization

4 identical tetrahedral sp3 orbitals

an sp3 orbital


Hybridization


• hybridized sp3 orbitals are unsymmetrical: one of two lobes is much larger, and form stronger bonds than unhybridized s or p orbitals

• unsymmetrical sp3 orbitals: two lobes of p-orbital have opposite signs

2s 2px

hybridization

an sp3 orbital

+_+ +_


H

CH H

H

109.5 o 110 pmbond angle(tetrahedral)

bond length

• C-H bond strength in methane: 438 kJ/mol (105 kcal/mol)

1.8 The Structure of Ethane• σ-overlap of two sp3 orbitals


C

H

HH

109.6 o

154 pm

C C

H

H

H

H

H

HCH3CH3

CC CC

C

H

HH

• C-H bond strength in ethane: 420 kJ/mol (100 kcal/mol)C-C bond strength in ethane: 376 kJ/mol (90 kcal/mol)

1.9 Hybridization: sp2 Orbitals and the Structure of Ethylene

• C2H4; double bond, planar


CH

H

121.7 o

133 pm

CH

H

C CH

H

H

HC C

H H

HH

116.6 o107.6 pm


s + 2 (p) → 3 (sp2)

3 identical sp2 orbitals

p

π-bond

σ-bond

• C-H bond strength in ethylene: 444 kJ/mol (106 kcal/mol)C=C bond strength in ethylene: 611 kJ/mol (146 kcal/mol)


Molecular orbital description C=C π-bond


combine

nodal plane

π-antibonding (π*) MO

π-bonding (π*) MO

Actually, p-orbitals have two lobes with opposite signs.

C OH

HC O

H

H

Lewis structure Line-bond structure


Practice

Line-bond structure, hybridization of carbon?

CH2O

sp2

C CC

H

HH

HH

H


Practice


CH3CH=CH2

sp3

sp2

1.10 Hybridization: sp Orbitals and the Structure of Acetylene

• C2H2; triple bond, linear


180 o

C CH HC CH H

106 pm120 pm


s + 1 (p) → 2 (sp)

2 identical sp orbitals

pp

π-bond

σ-bondπ-bond

• C-H bond strength in acetylene: 552 kJ/mol (132 kcal/mol)C≡C bond strength in acetylene: 835 kJ/mol (200 kcal/mol)



107.6106444C(sp2)-H

106132552C(sp)-H

110100420C(sp3)-H

120200835C(sp)≡C(sp)HC≡CH

133146611C(sp2)=C(sp2)H2C=CH2

15490376C(sp3)-C(sp3)CH3CH3

Bond Length(pm)kcal/molkJ/mol

BondMoleculeBond Strength

110105438C(sp3)-HCH4

Comparison of C/C and C/H bonds


Summary of Hybridizations

tetrahedral: sp3 hybrid

planar: sp2 hybrid

linear: sp hybrid C CH3C CH3

CH4 CH3-CH2-CH2-CH3

C CH3C CH3

H H

C CCH

HH

H


Practice


CH3C≡CH

sp3

sp

C C C

H

H

H

H


Practice Hybrid, Geometry?

CH2=C=CH2

C C C

H

H H

H

C CCH

H H

H

sp2 sp2

sp

C C C

H

H H

H

C C C

H

H H

H

1.11 Hybridization of Other Atoms: Nitrogen and Oxygen

• covalent bonds of other elements can also be described by hybridization


NH3 NH

H H

NH H

H

107.3 o

100.8 pm bond lengthN-H: 449 kJ/mol (107 kcal/mol)

sp3


H2O OH H

O

HH

104.5 o

95.8 pm bond length

O-H: 498 kJ/mol (119 kcal/mol)

sp3

C NHH

H


Practice

Line-bond structure, hybridizations?

Formaldimine, CH2NH

C NHH

H

sp2

sp2

OCH3

H


Practice Geometry?

CH3OHCH3NH3CCH3

CH3

N

CH3

CH3 PH

HH

PH3

1.12 Hybridization of Carbon Species

Carbocation


C C sp2

planar

CC

sp3

Carbanion

Chemical Toxicity and RiskChemistry @ Work

All chemicals are toxic to some extent, and the difference between help and harm is matter of degree.

4 x 10-4

1.73.210.62.41.5 x 10-2

17

Aflatoxin B1AspirinChloroformEthyl alcoholFormaldehydeSodium cyanideSodium cyclamate

LD50 (g/kg)Substance

LD50 value: the amount of a substance per kilogram body weight that is lethal to 50% of the test animals. The lower the value, the more toxic the substance.

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교과목명 유기화학(1) 학점 3 강의실 과B133 강의시간 수3,4 ......연습문제및출석및수업참여도: 50점 성적평가방법 강좌운영방식 교재중심강의