Introductory notes and review of general chemistry

CHE-310

Organic Chemistry I

Dr. James Lyle; office: NSM D-323

(310) 243-3388 or 243-3376

[email protected]

office hours: MWF: 9-10:00am & Tu: 8:00-9:00am

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 = 1004 exams @ 100 pts = 400final exam = 100homework required 600

Daily exams

No make ups! Drop two lowest scores. Begin at 10:00!

Daily Homework: Required!

(hold until called for)

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 99.985H2 1 1 0.015 C12 6 6 98.89C13 6 7 1.11C14 6 8 ...

Cl35 17 18 75.53Cl37 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: :NN:

HCN H:C:::N: H-CN:

.. ..H2CO H:C::O: H-C=O: .. | H H

atomic orbitals

s

p

detc.

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

B ABB

BA

B BB

B A B

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

..

..

HO

H

..

..

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:

1. Lewis dot structure

2. VSEPR hybridization shape of the molecule

3. dipoles for polar bonds

4. vector sum of the bond dipoles

5. vector sum = 0 non-polar molecule

6. vector sum 0 polar molecule

CO2 :O=C=O: sp linear

vector sum = 0non-polar molecule

H2O ..

H—O—H AX2E2 sp3 tetrahedral (bent) ..

H OH

vector sum 0

polar molecule!

HC

O HH

H

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



polar covalent => dipole-dipole &

Hydrogen bonding


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

ionic => ionic attractions

polar => dipole-dipole + H-bond

polar => dipole-dipole

non-polar => van der Waals

mp bp

350oC --

-94oC 65oC

-920C -21oC

-183oC –89oC

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/Baseshistoric:

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 weakeracid 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


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

H C C O HH H

H HH C O C HH H

H H

Documents

Introductory notes and review of general chemistry