Chemistry 343- Spring 2008

8/14/2019 Chemistry 343- Spring 2008

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Chemistry 343- Spring 2008Chapter 1- Carbon Compounds and Chemical Bonds

Basic Concepts of Structure

Valence= # of bonds an atom can typically form in a neutralcompound

Carbon = tetravalent; forms 4 bonds

Nitrogen = trivalent; forms 3 bondsOxygen = Divalent; forms 2 bonds, etc.

Molecular Formula (e.g.) C4H10 , C6H12O6

Isomerism= different compounds may have the same formulaIsomers: e.g., n-butanol and diethyl ether

CC

CC

OHH

H

H H

HHH

HH

CC

OC

CH H H H H

HH

H

H H

Structural (constitutional) isomer

Shape= geometry around a C atom depends on bondingarrangement4 single bonds tetrahedral2 single bonds, 1 double bond trigonal planar 1 single 1 triple OR 2 double bonds linear

Directly related to hybridization of C-atom orbitals


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Review of Some Relevant Principles From General ChemistryLewis dot structures: atoms, molecules, and ionsExamples:atoms in first row of periodic table

Li Be B C N O F Ne

molecules:

C H

H

HH O HH ClH C C

H

H

H

H

N H

H

HH O HH Cl C O

O

H

ions:

H

Rules: -show only valence e-- show any formal charge present- obey octet rule where appropriate

Formal Charge: a method of e- bookkeeping

NH3 + H+

NH4+

Atom Formal Charge = group # - net# of e- belonging to nucleus(lone pair counts as 2; shared pair contributes 1)

e.g., NH4+; Ns group # = 5; net # of e- = 4 formal charge = +1

For NH3; Ns group # = 5; net # of e-

= 5 formal charge = 0


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2. Octet Rule: An atom or ion is especially stable when its valenceshell has a noble gas configuration (i.e., 8 e- in most cases; 2e- for H, He, or Li+) some examples

C H

H

H

H H H F Li

Once the outer shell of any atom is filled, no more e- can be added

Larger atoms (3rd period and higher elements) may accommodatemore outer level e- due to their d and f orbitals, e.g., H2SO4, PCl5)

Species (atoms, ions, molecules) that do not fulfill the octet ruletend to be much less stable, and therefore much more reactive, e.gC H3+, BF3, Li .

3. Electronegativity: affinity of an atom for e-sFor Na ; electronegativity very low

For ; electronegativity very highCl Note extreme opposi positions in periodictable

Combination of electronegativity and octet rule suggests that : Na will lose 1 e- very readily to form Na+

will gain 1 e-

readily to give Cl-

Cl

Thus: Na + Na+ + Cl-ClStability isincreased !

4. Ionic Bonding: bonding based an electrostatic attraction

e.g., in solid NaCl:Cl-

Na+


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Each Na+ is equally attracted to 6 neighboring Cl- ionsEach Cl- is equally attracted to 6 neighboring Na+ ionsExplains physical and chemical properties (BP 1413oC; aqueoussolution, conducts electricity, etc.)

5. Covalent Bonding: give rise to very different structure types

Consider 2 Hydrogen atoms: H. H.Electronegativities are identicalE-shared to form a covalent bond (shared equally pure covalen

The octet rule is still obeyed for each atom

Consider CH4 (methane)

C,H electronegativties almost identicalC-H is considered a covalent bond

What if electronegativity difference is intermediate?polar covalent or polar bond

H H not H+

H-

C H

H

HH

e.g. H2O

2H . + O 2H + + O 2-

O HH


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The O-H bond is considered a polar bond

Consider their respective positions on the periodic table

Electronegativity increasing

Electron egativit yincrea sing

The electronegativity difference causes bonding e- to be polariz + and - are symbols used to depict partial + or - character

OH H

+ +Molecule has a net dipole momentand is considered polar

Other examples:

-

C OO

- -

+ -

H Cl

+

A polar molecule Has polar bonds, but dipolescancel; hence, not a polar molecule


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All organic molecules have covalent bonds of various types; someave ionic bonds too

Sometimes, atoms are bound by 2 or 3 covalent bonds

C CH

H

H

H

C OH

H

CH

CH

C OH

HThese havedouble bonds

6. Relevance of bonding type to chemical and physical properties

e.g. acetone

C

O

H3C CH 3

-

+

Dipoles can align to increase intermolecular attraction:influences BP, MP, other physical properties

O O + + O + - - -

Also, if areactive reagent comes along seeking a positive chargit may attack a + site and react; Conversely, if areactive reagentcomes along seeking a site of e- density, it may attack a - siteand reactWe will observe this phenomenom many, many times !

H 3C Br

- +

HO


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7. Resonance: A term used to describe a compound that cannot berepresented by a single Lewis dot structure

e.g. CH3 NO2 H 3 C NO

OH 3 C N

O

O

Experimentally:Both N-O bond lengths are same!Each of the resonance forms above are equally possible

CH3 NO2 behaves likeH3C NO

O12

12

-

-

Thus, CH3 NO2 is a resonance hybrid of these two forms shown

This is NOT an equilibrium!!

Symbol used to show that structures are related by resonanc

Note how one can show the conversion of one resonance form tothe other by showing only the movement of e - (curved arrows below)

H3C C

O

OH3C C

O

O

The use of this kind of arrow to show e- movement in a more

general sense (= curved arrow formalism) will be a critical toolused throughout this course!!!


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Molecular structures of covalently-bonded compoundsQ: How do we know them ?A: X-ray crystallography, spectral analysis

8. Molecular Geometry

A. Bond Length: Dependent on: Bond order (single, double, etc.) C=C vs. C-C Size of atoms C-C vs. H-H Electronegativity difference C-C vs. C-F

B. Bond Angles: Generally, angles are arranged so that e- pairs (loneor bonding) are as far apart as possible (VSEPR Theory )

If any combination of 4 atoms or lone pairs is associated with thecentral atom:

Shape is tetrahedral

C

H

HH

H109.5 o

NH

H

H

107 o

HO

H

105 o

Tetrahedral

Appears Pyramidal

Appears Bent


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Shape is trigonal planar

B

F

FF

120 o


Shape is linear C OO

180 o

All of this is based on

9. Orbitals and Hybridization

A. Background: Electrons possess characteristics of particles andwaves - - must be dealt with accordingly

Mathematical way of doing this involves Quantum Mechanics

Heisenberg Uncertainty Principle: We can only calculate the probability that an electron is in a certain region; hypothesize bounded areas for electron existence orbitals


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1Definitions:Atomic orbitals: mathematical wave functions ( ) describing energytates and corresponding regions in space in which a given e- can existn organic chemistry terms , an orbital is a region of space where the

probability of finding an electron is high

Orbitals are also known as energy levels two different orbitals wiorrespond to two distinct energy levels having certain discreet valu

Some Key Principles:

Aufbau Principle: lowest energy orbitals filled firstPauli Exclusion Principle: no more than two e- per orbital; spins mus

be paired (+1/2 and -1/2)Hunds Rule: for orbitals of equal energy, fill one e- at a time, only

add 2nd e- to an orbital once all have one unpaired e- each.

Review quantum numbers: must know 1s, 2s, and 2p orbitals and t

shapes (Figure 1.5, pg 21)

B. Simplest Case: the 1s orbital. Spherical shape, nucleus at center


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1- in a H atom occupy the 1s orbital- 90% probability that the e- iswithin 1 of nucleus; not uniform; probability goes up as r ecreases (r is the distance from the center of nucleus)

This e- can be depicted as a smear or cloud of e- density, or e-

robability about the nucleus

Wave function= ; magnitude decreases withincreasing r; mathematical sign can be + or -

Electron density= 2; magnitude decreases more

rapidly with increasing r; must always be +Mathematical sign is always + for 1s orbital(not true for any other)

1s orbital

C. 2s orbital: larger; also spherical, but is + beyond 1; -

inside 1 there is a spherical node at 1; 2

=0 at thatradius (e- cannot exist within that space)

e- can cross such nodes by virtue of their wavelike nature

Higher energy than 1s; avg. distance to nucleus is greater

2s orbital


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1The 2p orbitals: 3 of them ( p x , p y , p z )

Each has two lobes of equal size but opposite mathematical sign;

umbbell shape

Cannot occupy same space mutually perpendicular (2 p x , 2 p y , 2 p

Lobes of each are separated by a planar node going through nucleu

Higher energy than 1s and 2sThese (1s, 2s, and three 2p) are all the orbitals a carbon atom has !


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1E. Electron Configuration: for atoms of individual elements (Figure1.6, pg 22)

F. Molecular Orbitals (MOs): mathematical combinations of atomiorbitals ( s)

Rules:Can add or subtract AOs mathematically to get MOsIf we combine n orbitals, we must end up with n orbitalsOverall energy of new orbitals must be same after combination

Simplest Example: H2

Two 1s Atomic Orbitals (AOs) combine to form two new MO( and *)

One MO arises from addition of the AOs; the other fromsubtraction


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1Addition bonding MO (s);Subtraction antibonding MO (s*)

Orbital Diagram (Figure 1.10, pg 25)

Note that net overall energy of orbitals is the same as before, butthe system is more stable because the 2 previously unpaired e-scan both go into a lower energy orbital !Stability is gained through bond formation

A bond results from overlap of two s-orbitals = a -bondNote that Hunds rule and the Pauli Exclusion principle are still obe


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1Contrast this to what would happen if two He atoms came together

Two of the four e- would have to occupy the higher energy *antibonding orbital bond formation not favorable in thisinstance He exists in monoatomic form

Other Implications:

H2- can exist in theory, but would be unstable because the third e-would have to go into a * orbital

If enough energy were applied to an H2 molecule, an e- could beexcited to the * MO

In most common bonding situations, bonding MOs are filled,antibonding MOs are not


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1

G. Hybrid Atomic OrbitalsLets take a closer look at CH4

Experimentally, we know that:1. It has four equvalent C-H bonds2. It is tetrahedral

C

H

H H

H

Neither observation is consistent with the results expected fromcarbon AOs

Quantum Theory explains this:

The most favorable way for carbon to form bonds to four Hatoms is to mathematically combine 2s and three 2p orbitals to

give four new equivalent hybrid atomic orbitalscalled sp 3 orbitals (each is one parts to three partsp)sp 3 hybrid orbital energy is intermediate betweens and pvalence e- of carbon are redistributed- one e- per sp 3 orbitaleach sp 3 orbital can then combine with an H atoms 1s orbital toform MOs (only the bonding MOs are filled) form -

bonds

1s

2s

2p

MIX1s

sp 3

(x4)

A Free Carbon AtomThe Carbon atom of CH4


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1The shape of sp 3 orbitals is dictated by the math of the combination(Figure 1.11)

This same situation occurs anytime a C atom is bonded to four neighbors

e.g., look at ethane (Figure 1.17)

Note that free rotation can occur around the C-C bond (CH3s cantwirl around

Applies to more complex molecules as well


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1

Q : What about a situation where a C atom is only bonded to threeneighboring atoms, and shares two e- pairs with one of the neighbo

A: a difference in hybridization occurs!

.g. simplest case: ethene (a.k.a. ethylene, mol. Formula C2H4)C C

H

H

H

H

Experimentally:

Ethene has four equivalent C-H bonds; two to each C

Each C has trigonal planar geometry

All the atoms in the molecule are in a single plane

Molecule is rigid (no out of plane rotations or bending)

Neither unhybridized AOs nor sp 3 hybridization can explain this!


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1Quantum Theory states:

The most favorable way for carbon to form bonds with threeneighbors is to combine 2s and two of the three 2p orbitals to givthree equivalent hybrid atomic orbitals

called sp 2 orbitals (each is one part s to two parts p)

sp 2 energy is intermediate between s and p (lower Energy thansp 3)

Shape of sp 2 orbitals is dictated by the math of combination (Fig1.21)

Valence e- redistributed; one per sp 2; one in the remaining 2 p

Each sp 2 orbital can then combine with an H atoms 1s OR theneighboring carbonssp 2 orbital to form MOs

Only the bonding MOs are filled -bonds


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2The unhybridized p-orbitals can also overlap to form a (weaker)

bond; the 2nd bond of the double bond a -bond

p-orbitals must be parallel and inline with one another to overlapor bond formation to occur; explains rigidity; barrier to rotation

around C=C without breaking -bond !Fig 1.23, 1.24


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2This restricted rotation leads to a special type of isomerism:cis-trans isomerismC C

CH

H

H C

HC C

H

CH

H C

H

cis-2-butene trans-2-butene

cis-2-butene is a different compound from trans-2-butene !!These do not interconvert!

Cis and trans isomers are also known as geometrical isomers.This is one type of stereoisomerism (the connectivity is thesame, but the arrangement of atoms in space is somehow differen

Ok, but there are also molecules wherein a C atom is bound toonly two others!

e.g. simplest case: ethyne (a.k.a. acetylene)

C C HH


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2Experimentally:Ethyne has two equivalent C-H bonds; one to each CEach C has linear geometryall the atoms form a line

Neither unhybridized AOs,sp 2, nor sp 3 hybridization can explainthis !

Quantum Theory states:

The most favorable way for a C atom to form bonds with thesetwo neighbors is to combine the 2s and one of the three 2p orbitato give two equivalent hybrid atomic orbitals

called sp orbitals (each is one part s to one part p)

sp orbital energy intermediate between s and p (lower E thansp 2)

Shape of sp orbitals- again dictated by the math (Figure 1.27)


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2

Valence e- redistributed- 1 per sp orb.; 1 in each remaining 2p orb.

Each sp orbital can then combine with an H atoms 1s OR theneighboring carbonssp orbital to form MOs

Only the bonding MOs are filled -bonds

The two pairs of unhybridized p-orbitals can also overlap to formwo p-bonds; the other two bonds of the triple bond

Each pair of p-orbitals must again be parallel and inline for -bond

o formFigures 1.28, 1.29


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2Q: Are there any consequences of hybridization besides shape?

A: yes! Bond strengths and lengths

Double bonds are shorter and stronger than single bonds

Orbitals having more s-character form single bonds that arestronger and shorter (e.g., order of strength: sp-sp > sp-sp2 > sp-sp3 ~ sp2-sp2 > sp2-sp3 >sp3-sp3)

in general: more s-character lower energy

How does all this related back to VSEPR theory ?

Recall the shapes of ammonia and water; these can be

rationalized by hybridization!

Consider NH3: N bound only to 3 atoms, but there is also alone pair!

Molecule hybridizes to get the e- as far apart as possible!

Pyramid shape, but if lone pair is counted, it is tetrahedral!!

H2O does the same thing! It has two lone pairs and two bondsbent shape, but if you count lone pairs; it is ~ tetrahedral!!


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2This way of thinking applies to other molecules too.

Contrast NH3 with BF3:(NH3 has lone pair, BF3 does not; therefore NH3 has pyramidal shapeBF3 is flat (trigonal planar)

Another example: consider CO2(only two atoms around the central carbon, but each forms a doublbond with the carbon, hence the shape is linear)

H. Structural Formulas of Organic ChemistrySeveral different ways to draw organic molecules; Figure 1.36

Another common practice: lone pairs on O, N or other atoms aresometimes omitted, theyre just assumed to be there.


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2Condensed and bond-line (shorthand) presentations are especiallymportant; structures are complex!! Also, need to distinuish manysomers

Some structures are cyclic; easiest to draw as bond-line structuresC

C

CC

C

C

H HH

HH

HH H

HH

H

H~~

CC

CC

C

C

H

H

H

H

H

H

~~

Cyclohexane Benzene

In condensed formulas, its easier to draw structures with Cs and Habsent, known as implicit carbons and hydrogens

Also, will need a way to represent the three-dimensionality (3-D) othese structures; use wedges and dashes (Figure 1.37)

- (wedges) Depict atoms coming out of the plane of the pape- (dashes) Depicts atoms going into the plane of the paper

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Chemistry 343- Spring 2008

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